DESCRIPTIVE 
CHEMISTRY 


BY 


LYMAN   C.  NEWELL,  PH.D.  (JOHNS  HOPKINS) 

INSTRUCTOR   IN   CHEMISTRY,   STATE   NORMAL   SCHOOL,   LOWELL,    MASS. 
AUTHOR   OF   "EXPERIMENTAL   CHEMISTRY" 


BOSTON,  U.S.A. 

D.   C.   HEATH  &  CO.,   PUBLISHERS 
1903 


COPYRIGHT,  1903, 
BY   LYMAN    C.   NEWELL. 


ANTOIINE    LAURENT    LAVOISIER 

1743-1794 

THE    CELEBRATED    FRENCH    CHEMIST    WHO     LAID    THE     FOUNDATIONS    OF   CHEMISTRY 


PREFACE. 

THIS  book  is  intended  for  teachers  who  wish  to  emphasize  the  facts,  laws, 
theories,  and  applications  of  chemistry.  It  is  divided  into  two  parts.  Part  I 
contains  the  text,  together  with  exercises  and  problems.  Part  II  contains  the 
experiments.  The  text  has  been  selected  and  arranged  with  special  refer- 
ence to  the  needs  of  teachers  as  well  as  to  the  capacity  of  students.  The 
experiments  have  been  prepared  to  meet  the  needs  of  those  schools  in  which 
the  laboratory  facilities  are  limited  or  the  time  for  chemistry  is  short. 

The  point  of  view  differs  from  that  in  the  author's  "  Experimental  Chem- 
istry," but  the  spirit  is  the  same.  The  two  books  are  companion  volumes, 
though  of  course  they'  can  be  used  independently.  The  cordial  reception 
given  the  "  Experimental  Chemistry "  shows  that  many  teachers  are  empha- 
sizing the  experimental  side  of  chemistry.  These  teachers  will  find  Part  I 
of  the  "  Descriptive  Chemistry "  a  serviceable  companion  book  both  in  the 
laboratory  and  class  room.  It  has  been  bound  as  a  separate  volume  to  meet 
such  a  use. 

Solutions  of  problems,  answers  to  some  of  the  exercises,  and  references  to 
the  literature  have  been  put  in  a  separate  Teacher's  Handbook. 

The  manuscript  has  been  read  by  Dr.  William  B.  Schober,  Lehigh  Uni- 
versity, Bethlehem,  Pennsylvania;  Mr.  Franklin  T.  Kurt,  Chauncey  Hall 
School,  Boston,  Massachusetts;  and  Mr.  George  M.  Turner,  Masten  Park 
High  School,  Buffalo,  New  York.  The  chapters  on  theory  were  also  read  by 
Dr.  Alexander  Smith  of  the  University  of  Chicago,  and  the  chapters  on 
carbon  by  Dr.  James  F.  Norris  of  the  Massachusetts  Institute  of  Technology. 
The  proof  has  been  read  by  Dr.  E.  H.  Kraus,  High  School,  Syracuse,  New 
York;  Professor  E.  S.  Babcock,  Alfred  University,  Alfred,  New  York;  and  Mr. 
E.  R.  Whitney,  High  School,  Binghamton,  New  York.  The  author  is  grateful 
to  these  teachers  for  their  criticism,  but  he  assumes  all  responsibility  for  any 
errors  which  may  be  detected. 

L.  C.  N. 

LOWELL,  MASS., 

239204 

iii 


„ 


,... 


CONTENTS. 

PART   I.  ,  f  } 

CHAPTER  PAGE 

-I.    PHYSICAL  AND  CHEMICAL  CHANGES  —  CHEMICAL  ACTION,  ~r 

CHEMICAL  ENERGY  —  ELEMENTS  —  COMPOUNDS      .       '.        i 
II.    OXYGEN  —  LAWS  OF  CHARLES  AND  BOYLE  —  OZONE     .        .      1 1 

III.  HYDROGEN       .        .        .        ....        .        .        .        .23 

IV.  GENERAL  PROPERTIES  OF  WATER 31 

V.    COMPOSITION  OF  WATER  —  HYDROGEN  DIOXIDE   ...      50 

fVI.    THE  ATMOSPHERE  —  NITROGEN 61 

VII.  LAW  AND  THEORY  —  LAWS  OF  DEFINITE  AND  MULTI-J^ 
PROPORTIONS  —  ATOMIC  THEORY  —  ATOMS  AND  MOLE- 
CULES —  SYMBOLS  AND  FORMULAS  —  EQUATIONS  .  .  75 

VIII.    ACIDS,  BASES,  AND  SALTS 87 

IX.    EQUIVALENTS — ATOMIC  AND  MOLECULAR  WEIGHTS  —  CHEMI- 
CAL  CALCULATIONS  —  QUANTITATIVE   SIGNIFICANCE   OF 
EQUATIONS        ....        -^^^^^-        •        •     100 
X.    LIGHT,  HEAT,  ELECTRICITY,  AND  CHEMICAJMP»N     .        .    in 
XI.    CHLORINE  AND  HYDROCHLORIC  ACID     .        .        .        .        .133 
—  XII.    AMMONIA  —  NITRIC  ACID  AND  NITRATES  —  AQUA  REGIA  — 

^p        OXIDES  OF  NITROGEN ^  14*1 

XIII.  GASES  —  GAY-LUSSAC'S  LAW  —  AVOGADRO'S    HYPOTHESIS  — 

VAPOR  DENSITY  --(M.OLECULAR  AND  ATOMIC  WEIGHTS 
—  MOLECULAR  FORMULA  —  MOLECULAR  EQUATIONS  — 
VALENCE 166 

XIV.  CARBON  AND  ITS  OXIDES  —  CYANOGEN 181 

XV.    MF.THANE  —  ETHYLENE  —  ACETYLENE  —  ILLUMINATING  GAS 

^.  —  FLAME  —  BUNSEN  BURNER  —  OXIDIZING  AND  REDUC- 
ING FLAME 202 


J  Contents. 

HAPTER  PAGE 

XVI.  FLUORINE  —  BROMINE  —  IODINE  .        .        .       .        .        .    225 

^+-  XVII.    SULPHUR  AND  ITS  COMPOUNDS 235 

XVIII.    SILICON  —  BORON .255 

XIX.  PHOSPHORUS  —  ARSENIC  —  ANTIMONY  —  BISMUTH      .        .    265 

XX.    METALS .278 

XXI.    SODIUM  —  POTASSIUM  —  LITHIUM 284 

XXII.  COPPER  —  SILVER  —  GOLD    .        .        .        .                .        .301 

XXIII.  CALCIUM  —  STRONTIUM  —  BARIUM 319 

XXIV.  MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY      .        .        .331 
XXV.    ALUMINIUM  .        . 343 

XXVI.    TIN  — LEAD .        .        .        .354 

XXVII.    CHROMIUM  —  MANGANESE 365 

kxVIII.    IRON  —  NICKEL  —  COBALT   . 373 

\XXIX.  PLATINUM  AND  ASSOCIATED  METALS   .        .        «..'..«        •    392 

\>J  XXX.  PERIODIC  LAW  —  SPECTRUM  ANALYSIS         ....    396 

4  ,_AXI.    SOME  COMMON  ORGANIC  COMPOUNDS 405 

APPENDIX  ............    437 


PART    I 

DESCRIPTIVE 
CHEMISTRY 


DESCRIPTIVE  CHEMISTRY. 


CHAPTER    I. 
INTRODUCTION. 

CHEMISTRY  is  a  branch  of  natural  science.  It  deals 
with  the  properties  of  matter,  the  changes  which  affect 
the  composition  of  matter,  with  numerous  laws  and 
theories,  and  with  the  manufacture  of  a  vast  number  of 
different  substances  indispensable  to  the  welfare  of  man- 
kind. 

Properties  of  Matter.  —  Different  substances  are  recog- 
nized and  distinguished  by  their  properties.  Color,  odor, 
taste,  weight,  and  solubility  are  familiar  properties ;  but  to 
these  must  be  added  behavior  with  heat,  light,  and  electric- 
ity, and  especially  the  action  of  different  kinds  of  matter 
upon  each  other. 

Physical  and  Chemical  Changes.  —  Observation  shows 
that  the  properties  of  matter  can  be  changed.  Sometimes 
the  change  is  only  temporary,  as  in  the  freezing  of  water, 
or  in  the  melting  of  iron.  Such  changes  are  called  physi- 
cal changes.  But  often  the  change  is  permanent,  as  in 
the  burning  of  coal,  or  the  digestion  of  food.  Such 
changes  are  called  chemical  changes.  In  physical 
changes  the  original  properties  reappear  after  the  cause 
of  the  change  has  been  removed.  But  chemical  changes 


2  Descriptive  ^Chemistry. 

affect  the  essential  nature  of  a  substance.  They  are 
fundamental.  Removal  of  the  cause  of  a  chemical  change 
does  not  restore  the  original  properties  of  the  substance. 
Thus,  coal  is  readily  changed  into  ashes  and  invisible 
gases,  but  the  ashes  and  gases  do  not  reunite  into  coal 
after  the  heat  has  been  removed.  Another  essential  char- 
acteristic of  chemical  changes  is  the  formation  of  one  or 
more  kinds  of  matter  different  from  the  original  substance. 
Thus,  water  may  be  decomposed  by  electricity  into  two 
gases  —  hydrogen  and  oxygen.  This  is  a  chemical  change, 
because  (i)  the  water  has  disappeared,  its  identity  is  lost, 
it  has  been  permanently  changed,  and  (2)  other  kinds  of 
matter  have  been  formed,  which  are  totally  unlike  water. 
Chemistry  is  largely  a  study  of  chemical  changes. 

The  different  changes  which  matter  undergoes  furnish 
a  convenient  basis  for  the  classification  of  properties. 
Thus,  we  call  physical  properties  those  which  accompany 
physical  changes ;  while  chemical  properties  require  a 
chemical  change  for  their  manifestation.  Thus,  the  color, 
luster,  specific  gravity,  melting  point,  and  capacity  to  con- 
duct electricity  are  physical  properties  of  copper;  but  it 
displays  chemical  properties  when  it  is  heated,  or  when 
acted  upon  by  acids,  sulphur,  and  other  substances. 

Examples  of  simple  physical  changes  are  the  formation  of  ice  or  steam 
from  water,  the  electrification  of  a  copper  trolley  wire,  the  production  of 
colors  in  the  sky,  the  magnetization  of  iron  in  a  dynamo  or  magnet,  and 
\the  melting  of  iron  in  a  foundry.  Familiar  chemical  changes  are  the 
rusting  of  iron,  the  growth  of  plants,  the  burning  of  oil  in  a  lamp,  the 
decay  of  fruit,  and  the  souring  of  milk. 

Chemical  changes  are  often  complex.  In  many  in- 
stances they  are  caused  by  heat,  and  usually  they  produce 
heat.  In  general,  the  velocity  of  chemical  change  in- 
creases with  rise  of  temperature.  Light  induces  chemical 


Introduction.  3 

changes,  as  in  growing  plants  and  on  photographic  plates. 
Electricity  is  involved  in  many  chemical  changes,  a  vast 
industry  having  recently  grown  up  in  this  field.  Contact 
is  necessary  for  chemical  change,  and  many  substances 
must  be  pressed  together,  intimately  mixed,  or  dissolved 
before  they  will  interact. 

Physical  and  chemical  changes  are  closely  related. 
They  usually  accompany  each  other,  and  are  often  insep- 
arable. If  the  essential  change  in  a  substance  or  sub- 
stances is  chemical,  then  the  substances  are  said  to 
undergo  chemical  action.  Very  often  the  chemical  action 
involves  several  substances.  The  substances  are  then  said 
to  interact  or  react,  and  the  series  of  changes  is  called  a 
reaction.  Thus,  when  zinc  is  added  to  nitric  acid,  the 
chemical  action  which  occurs  is  manifested  by  the  forma- 
tion of  a  brown  gas  and  the  disappearance  of  the  zinc. 
The  zinc  and  acid  interact,  and  tlie  chemical  changes  can 
be  classified  as  due  to  the  reaction  between  zinc  and  nitric 
acid. 

Classes  of  Chemical  Action.  —  There  are  four  general 
kinds  of  chemical  action,  (i)  Analysis  or  decomposition 
is  the  separation  of  matter  into  its  components.  Thus, 
heat  decomposes  wood,  and  the  juices  of  our  bodies  de- 
compose food.  (2)  Synthesis  or  combination  is  the  union 
of  different  kinds,  or  sometimes  the  same  kind,  of  matter. 
For  example,  the  gases,  hydrogen  and  oxygen,  may  be 
made  to  unite  and  form  water  by  passing  an  electric  spark 
through  them.  (3)  Substitution  is  the  replacement  of 
one  kind  of  matter  by  another.  When  zinc  is  added  to 
hydrochloric  acid,  the  hydrogen  leaves  the  acid,  and  zinc 
takes  its  place.  (4)  Sometimes  parts  of  different  sub- 
stances exchange  places ;  this  kind  of  change  is  called 
metathesis  or  double  decomposition.  If  silver  nitrate  is 


4  Descriptive  Chemistry. 

added  to  hydrochloric  acid,  the  silver  and  hydrogen  ex- 
change places,  forming  silver  chloride  and  nitric  acid. 
These  four  kinds  of  chemical  changes  will  be  fully  illus- 
trated and  studied  in  the  succeeding  pages. 

Chemical  Energy.  —  We  learn  in  physics  that  heat, 
light,  and  electricity  are  different  forms  of  energy.  They 
produce  special  changes.  It  is  also  possible  to  transform 
the  different  kinds  of  energy  into  each  other.  Thus,  elec- 
tricity is  generated  from  the  heat  liberated  by  burning 
coal,  and  electricity  in  turn  may  be  transformed  into  light. 
In  chemistry  we  study  another  kind  of  energy,  called 
chemical  energy,  chemical  attraction,  or  chemism.  This 
is  the  immediate  agent  involved  in  chemical  change.  Com- 
bination and  decomposition  are  due  to  its  operation. 
Chemical  energy  may  be  transformed  into  light,  electricity, 
and  heat,  and  vice  versa.  Appreciable  heat  often  accom- 
panies chemical  changes,  and  we  shall  have  many  illustra- 
tions of  the  intimate  relation  between  heat  and  chemical 
energy.  Electricity  is  produced  in  an  electric  battery  by 
chemical  action.  Light  is  one  result  of  the  chemical 
action  called  combustion  or  burning.  In  fact,  every  chemi- 
cal change  is  accompanied  by  an  energy  change  of  some 
kind,  and  in  such  transformations  all  the  energy  can  be 
accounted  for,  none  is  lost  or  gained. 

Chemical  energy  is  an  essential  factor  in  all  chemical 
changes,  but  we  know  little  or  nothing  of  its  nature.  We 
can  only  study  its  results  and  its  manner  of  action. 

Conservation  of  Matter.  —  In  chemical  changes  matter 
is  not  created  or  destroyed.  It  is  often  transformed,  and 
apparently  lost,  but  the  total  weight  of  the  substances  par- 
ticipating in  any  chemical  change  is  always  the  same. 
The  fact  that  matter  is  indestructible  was  first  demon- 


Introduction.  5 

strated  by  the  French  chemist,  Lavoisier  (1743-1794),  and 
countless  observers  have  since  shown  that  it  is  a  funda- 
mental law  of  chemistry.  The  law  is  called  the  Law  of 
the  Conservation  of  Matter,  and  is  often  stated  thus :  - 

No  weight  is  lost  or  gained  in  a  chemical  change. 

Chemical  Elements.  —  Study  of  the  constitution  of 
matter  shows  that  some  kinds  can  be  decomposed  into 
substances  totally  unlike  the  original  matter.  Water,  for 
example,  is  easily  decomposed  into  the  gases,  hydrogen 
and  oxygen,  which  are  entirely  different  from  water.  But 
it  is  impossible  by  any  known  process  to  obtain  from 
some  kinds  of  matter  substances  which  have  simpler  prop- 
erties than  the  original  substance.  Thus,  neither  oxygen 
nor  hydrogen  can  be  decomposed  by  any  known  means. 
Iron  and  the  familiar  metals  likewise  cannot  be  divided 
chemically  into  two  or  more  substances,  nor  can  they  be 
transformed  into  each  other.  They  are  fundamental  sub- 
stances. We  can  add  other  substances  to  them,  but  we 
cannot  get  simpler  substances  from  them,  nor  can  we 
transform  them  into  simpler  substances.  Iron  contains 
nothing  but  iron.  The  substances  which  have  such  simple 
properties  and  at  present  defy  decomposition  and  trans- 
formation are  called  the  chemical  elements.  They  are 
analogous  to  the  letters  of  the  alphabet,  and  by  their  vari- 
ous combinations  make  up  the  matter  of  the  universe,  some- 
what as  letters  form  words. 

There  are  about  eighty  elements.  Probably  there  are 
some  undiscovered,  but  it  is  generally  believed  that  the 
present  number  will  not  be  largely  increased. 

Each  element  is  designated  by  a  symbol,  which  is  an 
abbreviation  of  its  name.  The  following  is  an  alphabeti- 
cal— 


Descriptive  Chemistry. 
TABLE  OF  THE  IMPORTANT  ELEMENTS. 


NAME. 

SYMBOL. 

NAME. 

SYMBOL. 

Aluminium     .... 

Al 

Lead  .... 

Pb 

Antimony       .... 

Sb 

Lithium  

Li 

Arsenic      
Barium 

As 
Ba 

Magnesium  .... 
Manganese 

Mg 
Mn 

Bismuth    

Bi 

Mercury  

Hg 

Boron  

B 

Nickel 

Ni 

Bromine    
Cadmium 

Br 

Cd 

Nitrogen       .... 
Oxygen 

N 

o 

Calcium 

Ca 

Phosphorus 

p 

Carbon      

c 

Platinum 

Pt 

Chlorine    

Cl 

Potassium 

K 

Chromium      .... 
Cobalt                       .     . 

Cr 
Co 

Silicon     
Silver 

Si 
Ae1 

Copper       

Cu 

Sodium 

"8 

Na 

Fluorine    
Gold 

F 
Au 

Strontium     .... 
Sulphur 

Sr 

s 

Hydrogen 

H 

Tin 

Sn 

Iodine  

I 

Zinc    

Zn 

Iron 

Fe 

Of  the  above  elements  only  eight  are  abundant  in  the 
earth's  crust,  as  may  be  seen  by  a  — 

TABLE  OF  THE  APPROXIMATE  COMPOSITION  OF  THE  EARTH'S  CRUST 
(BY  WEIGHT). 


ELEMENT. 


Oxygen 
Silicon 
Aluminium 
Iron    . 
Calcium 
Magnesium 
Potassium  . 
Sodium 


Total 


PER  CENT. 

47.29 

27.21 
7.8l 
5.46 

3-77 
2.68 
2.40 
2.36 
98.98 


Introduction. 


The  atmosphere  contains  about  20  per  cent  of  oxygen 
and  79  per  cent  of  nitrogen  in  the  free  state.  The  ocean 
contains  about  86  per  cent  of  oxygen,  1 1  per  cent  of  hydro- 
gen, and  2  per  cent  of  chlorine  in  combined  states.  It  is 
clear  that  the  globe,  as  we  know  it,  is  made  up  of  a  very 
few  elements. 

Many  of  the  familiar  metals  are  elements,  e.g.  lead,  zinc, 
tin,  copper,  iron,  gold,  and  silver.  Other  elements  besides 
the  metals  are  solids,  such  as  sulphur,  carbon,  and  phos- 
phorus ;  two  are  liquid,  viz.  bromine  and  mercury ;  while 
several  are  the  common  gases,  oxygen,  nitrogen,  and  hydro- 
gen. Many  are  important  simply  because  they  are  com- 
bined with  other  elements,  especially  silicon,  which  is 
found  in  most  rocks,  and  calcium,  which  is  a  component 
of  limestone. 

The  following  is  a  — 

TABLE  OF  THE  UNCOMMON  ELEMENTS. 


NAME. 

SYMBOL. 

NAME. 

SYMBOL. 

Ar^on  

A 

Prasedymium 

Pr 

Beryllium 

Be 

Rhodium 

Rh 

Caesium 

Cs 

Rubidium 

Rb 

Cerium      
Erbium 

Ce 
Er 

Ruthenium      .... 
Samarium 

Ru 

Sm 

Gallium          .... 

Ga 

Scandium             . 

Sc 

Germanium 

Ge 

Selenium 

Se 

Glucinum 

Gl 

Tantalum 

Ta 

Helium      

He 

Tellurium      »  .     .     .     . 

Te 

Indium 

In 

Thallium 

Tl 

Iridium 

Ir 

Thorium 

Th 

Krypton    

Kr 

Titanium    

Ti 

Lanthanum 

La 

Tungsten 

W 

Molybdenum  

Mo 

Uranium     

u 

Neodymium 

Nd 

Vanadium             •          . 

v 

Neon    . 

Ne 

Xenon         

Xe 

Niobium    

Nb 

Yb 

Osmium 

Os 

Yttrium       

Yt 

Palladium 

Pd 

Zirconium        .... 

Zr 

8 


Descriptive  Chemistry. 


Chemical  Symbols  are  usually  the  first  letter  of  the 
name  of  the  element.  Thus,  O  is  the  symbol  of  oxygen, 
H  of  hydrogen,  N  of  nitrogen.  Since  several  elements 
have  the  same  initial  letter,  the  symbol  of  some  elements 
contains  two  letters.  Thus,  C  represents  carbon,  while 
the  symbol  of  calcium  is  Ca,  of  chlorine  Cl,  of  chromium 
Cr,  and  of  copper  Cu.  The  symbols  of  several  elements, 
especially  the  metals  so  long  known,  are  derived  from 
their  Latin  names,  as  may  be  seen  from  a  — 


TABLE  OF  LATIN  SYMBOLS. 


ELEMENT. 

LATIN  NAME. 

SYMBOL. 

ELEMENT. 

LATIN  NAME. 

SYMBOL. 

Antimony 

Stibium 

Sb 

Mercury 

Hydrargyrum 

Hg 

Copper 

Cuprum 

Cu 

Potassium 

Kalium 

K 

Gold 

Aurum 

Au 

Silver 

Argentum 

Ag 

Iron 

Ferrum 

Fe 

Sodium 

Natrium 

Na 

Lead 

Plumbum 

Pb 

Tin 

Stannum 

Sn 

Symbols  always  begin  with  a  capital,  and  are  not  followed  by  a 
period.  They  should  be  learned  by  actual  use.  Their  significance 
will  be  explained  in  later  chapters. 

Chemical  Compounds.  —  When  elements  unite  with 
each  other  the  product  of  the  union  is  a  chemical  com- 
pound. The  elements  which  make  up  a  chemical  com- 
pound are  called  components.  Chemical  compounds  have 
three  essential  characteristics,  (i)  Their  components  are 
held  together  by  chemical  attraction.  The  hydrogen  and 
oxygen,  which  are  the  components  of  water,  cannot  be 
separated  unless  their  attraction  for  each  other  is  over- 
come by  heat,  electricity,  or  some  other  agent.  (2)  In  any 
given  chemical  compound  the  components  are  always  in 


Introduction.  g 

the  same  ratio.  Thus,  pure  common  salt,  however  pre- 
pared or  wherever  found,  always  contains  39.32  per  cent 
of  sodium  and  60.68  per  cent  of  chlorine.  So  also  water 
always  contains  eight  parts  (by  weight)  of  oxygen  and  one 
of  hydrogen.  Facts  similar  to  these  might  be  given  cover- 
ing all  cases  examined.  Such  facts  illustrate  the  general 
principle  that  chemical  action  proceeds  according  to  laws. 
(3)  In  chemical  compounds  the  identity  of  the  components 
is  lost.  Thus,  the  red  metal,  copper,  the  yellow  solid, 
sulphur,  and  the  invisible  gas,  oxygen,  are  the  components 
of  the  blue  solid,  copper  sulphate. 

Chemical  compounds  must  not  be  confused  with  mixtures.  The 
parts  of  a  mixture  may  vary  in  nature  and  in  proportion ;  they  are  also 
held  together  loosely,  and  may  often  be  separated  by  some  mechanical 
operation,  as  filtering  or  sifting.  A  mixture,  too,  often  has  properties 
similar  to  its  parts. 

EXERCISES.1 

1.  State  three  properties  of  (a)  glass,  (<£)  wood,  (c}  water,  (W)  paper, 
(e)  air. 

2.  Give  three  illustrations  of  (a)  physical  changes  and  (6)  chemical 
changes  occurring  in  everyday  life. 

3.  Are  the  following  changes  physical  or  chemical?     (a}  Burning 
of  wood,  (£)  melting  of  butter,  (c)  freezing  an  ice-cream  mixture,  (d} 
weathering  (i.e.  decay)  of  granite,   (e)  tarnishing  of  brass  and  other 
metals,  (/)  formation  of  snow,  (g)  developing  a  photographic  plate,  (h) 
seasoning  of  wood,  (/)  formation  of  dew,  (/)  disappearance  of  a  fog. 

4.  What  ai^ls  and  what  retards  chemical  change?     What  often  ac- 
companies it? 

5.  What  physical  change  accompanies  (a}  the  burning  of  coal,  (6) 
the  action  of  an  electric  battery,  (c)  the  burning  of  a  match  ? 

6.  Give  an  illustration  of  the  transformation  of  chemical  energy  into 
heat,  light,  or  electricity. 

7.  State  the  law  of  the  conservation  of  matter. 

1  These  exercises  are  intended  for  review  work. 


io  Descriptive  Chemistry. 

8.  (rt)  Name  five  elements  with  which  you  are  familiar.     (£)  Name 
the  eight  most  abundant  elements  in  the  earth's  crust  in  their  order. 

9.  What  common  metals  are  elements? 

10.  How  do  elements  and  compounds  essentially  differ?     Could  you 
prepare  (a)  a  compound  from  elements,  (^)  elements  from  a  compound, 
and  (c)  elements  from  elements? 

11.  Define  (a)  chemistry,  (£)  physical  change,  (c)  chemical  change, 
(d)  chemical  action,  (^)  analysis,  (/")  synthesis,  (g)  metathesis,  (^)  sub- 
stitution, (/)  element,  (/)  compound,  (£)  mixture,  (/)  symbol. 

12.  Review  or  learn  the  metric  system  (see  Appendix,  §  i). 


PROBLEMS. 

Perform  the  problems  in  the  Appendix,  §  i, 


CHAPTER    II. 
OXYGEN. 

OXYGEN  has  played  an  important  part  in  the  develop- 
ment of  chemistry,  and  is  an  appropriate  element  with 
which  to  begin  a  systematic  study  of  this  science. 

Occurrence.  —  Oxygen  is  the  most  abundant  atid  widely 
distributed  of  the  elements.  Mixed  with  nitrogen  and  a 
few  other  gases,  it  forms  one  fifth  (by  volume)  of  the 
atmosphere.  Combined  with  hydrogen,  it  constitutes 
eight  ninths  (by  weight)  of  water;  combined  with  silicon 
and  certain  metals,  it  makes  up  nearly  half  of  the  earth's 
crust;  while  compounds  of  oxygen,  carbon,  and  hydrogen 
form  a  large  part  of  animal  and  vegetable  matter.  Starch, 
for  example,  which  is  a  constituent  of  all  plants,  contains 
about  50  per  cent  oxygen. 

Preparation.  —  Oxygen  may  be  prepared  from  its  com- 
pounds or  from  air.  It  was  first  prepared  by  decomposing 
a  red  compound  of  oxygen  and  mercury.  When  heated 
in  a  hard  glass  tube,  this  compound  decomposes  into 
oxygen  and  mercury ;  the  oxygen  is  collected  over  water 
in  a  pneumatic  trough,  and  the  mercury  condenses  as 
globules  or  a  film  on  the  upper  part  of  the  tube.  This 
experiment  is  historically  interesting,  because  it  was  first 
performed  by  Priestley,  the  discoverer  of  oxygen. 

The  gas  is  often  prepared  by  decomposing  potassium 
chlorate  —  a  compound  of  oxygen,  chlorine,  and  potassium. 
Heated  to  a  rather  high  temperature,  the  potassium  chlo- 


12  Descriptive  Chemistry. 

rate  passes  through  a  series  of  changes ;  as  a  final  result, 
the  oxygen  is  set  free,  and  potassium  chloride,  a  white 
solid,  remains  behind. 

Oxygen  is  most  conveniently  prepared  by  heating  a 
mixture  of  potassium  chlorate  and  manganese  dioxide  in  a 
glass  or  metal  vessel.  The  gas  is  liberated  freely  from 
this  mixture  at  a  lower  temperature  than  when  either 
compound  is  heated  alone. 

The  manganese  dioxide  may  be  recovered  unchanged  at  the  close  of 
the  experiment.  It  takes  some  part  in  the  chemical  changes,  but  just 
what  is  not  definitely  known.  It  has  been  suggested  that  the  manganese 
dioxide  combines  at  first  with  oxygen,  thereby  forming  another  coin- 
pound  of  manganese  richer  in  oxygen  than  the  dioxide,  but  so  unstable 
that  when  heated  it  yields  oxygen  and  manganese  dioxide. 

Large  quantities  of  oxygen  may  be  prepared  by  heating  a  mixture  of 
potassium  chlorate  and  manganese  dioxide  in  a  copper  or  iron  retort. 
Other  commercial  processes  are  used.  In  Erin's  process,  which  is  oper- 
ated largely  in  England,  purified  air  is  forced  by  a  pump  over  barium 
oxide  heated  to  700°  C.,1  thereby  forming  barium  dioxide.  The  air  sup- 
ply is  then  cut  off,  and  the  pressure  in  the  retorts  reduced  by  reversing 
the  pump.  This  operation  changes  the  barium  dioxide  into  barium  oxide 
and  oxygen.  The  gas  is  drawn  off  into  a  reservoir.  The  process  is 
then  repeated.  A  kilogram  of  barium  oxide  yields  about  ten  liters  of 
oxygen  at  a  single  operation.2 

Oxygen  can  be  prepared  from  liquid  air  (see  Liquid  Air).  By  evapo- 
ration at  the  ordinary  temperature  and  pressure,  the  nitrogen  escapes 
from  the  liquid  air  more  rapidly  than  the  oxygen,  leaving  finally  a  liquid 
which  is  nearly  pure  oxygen.  Unlimited  quantities  of  oxygen  may  thus 
be  cheaply  prepared  from  the  air.  This  method  awaits  development. 

Properties.  —  Oxygen  gas  has  no  color,  odor,  or  taste. 
It  is  slightly  heavier  than  air.  It  is  somewhat  soluble  in 

1  C.  is  the  abbreviation  of  "  centigrade,"  which  is  the  name  of  the  thermometer 
used  in  science.    According  to  this  thermometer  water  boils  at  100°  and  freezes  at 
o°  (see  Appendix,  §  2). 

2  "  Kilogram  "  and  "  liter"  are  denominations  of  the  Metric  System  of  Weights 
and  Measures.    This  system  should  be  learned  or  reviewed  (see  Appendix,  §  i). 


Oxygen.  13 

water,  but  the  presence  of  even  a.  small  proportion  in 
water  is  exceedingly  important.  Fish  die  in  water  con- 
taining no  oxygen;  and  the  oxygen  absorbed  by  flowing 
water  helps  keep  it  free  from  organic  matter.  (See  Decay, 
below.) 

The  density  of  oxygen  gas  is  1.105  (air  =  i).  One  hundred  liters 
of  water  dissolve  only  about  three  liters  of  oxygen  under  ordinary 
conditions. 

The  chemical  activity  of  oxygen  is  its  most  striking 
property.  It  combines  with  all  the  other  elements  except 
fluorine,  bromine,  and  the  inert  gases  recently  discovered 
in  the  atmosphere.  With  most  of  them  the  union  is 
direct,  and  is  often  accompanied  by  light  and  heat, 
though  the  temperature  at  which  combination  occurs 
varies  between  wide  limits.  At  the  ordinary  temperature 
it  unites  with  phosphorus,  as  may  be  seen  by  the  glow  and 
fumes  when  the  end  of  a  match  is  rubbed,  especially  in  a 
dark  room.  Metals,  such  as  iron,  lead,  zinc,  and  copper, 
tarnish  or  rust  easily,  i.e.  they  combine  with  the  oxygen 
of  the  air.  The  chemical  activity  of  oxygen  at  high  tem- 
peratures is  readily  shown  by  putting  burning  substances 
into  it.  All  burn  vividly  in  oxygen. 

When  a  glowing  stick  of  wood  is  put  into  oxygen,  the  stick  instantly 
bursts  into  a  flame ;  and  if  left  in'  the  oxygen,  the  wood  continues  to 
burn  brightly  until  the  gas  is  exhausted.  If  glowing  charcoal  is  put 
into  oxygen,  the  charcoal  burns  violently,  and  throws  off  showers  of 
sparks.  Sulphur  burns  in  air  with  a  small,  blue  flame,  but  in  oxygen 
the  flame  is  much  larger  and  brighter.  The  flame  in  both  cases  is 
accompanied  by  fumes  which  smell  like  a  burning  sulphur  match.  Iron 
wire  does  not  burn  in  air,  but  if  the  end  is  coated  with  burning  sulphur 
and  then  put  into  oxygen,  the  wire  burns  vividly,  throwing  off  a  shower 
of  sparks  ;  when  the  flame  has  disappeared,  a  globule  of  red-hot  iron  is 
often  seen  on  the  end  of  the  wire ;  and  sometimes  the  inside  of  the 
bottle  is  coated  with  a  reddish  powder,  which  is  mainly  a  compound 


14  Descriptive  Chemistry. 

of  iron  and  oxygen.  Iron  and  oxygen  combine  at  a  higher  tempera- 
ture than  do  sulphur  and  oxygen,  so  sulphur  is  used  to  set  fire  to  the 
iron.  On  the  other  hand,  if  lighted  magnesium  is  put  into  oxygen,  the 
burning  metal  instantly  becomes  surrounded  with  a  dazzling  flame,  and 
burns  rapidly  to  a  white  powder,  thus  showing  that  the  temperature  at 
which  it  combines  with  oxygen  is  much  lower  than  that  required  by  iron. 

Oxidation.  —  When  sulphur,  iron,  magnesium,  and  car- 
bon (in  wood  and  charcoal),  and  other  elements  burn  in 
oxygen,  they  combine  with  it.  This  chemical  change  is 
called  oxidation. 

The  fact  that  oxidation  is  merely  a  combining  with  oxygen  may  be 
easily  verified.  It  has  been  repeatedly  shown  that  oxygen  is  one  con- 
stituent of  all  the  products  formed  by  burning  substances  in  that  gas. 
Thus,  carbon  forms  an  invisible  gas  called  carbon  dioxide,  which  is  a 
compound  of  carbon  and  oxygen.  Similarly,  sulphur,  iron,  and  magne- 
sium form  compounds  of  these  elements  and  oxygen.  These  facts  may 
be  further  verified  by  a  simple  experiment.  If  mercury  is  heated,  it 
gains  in  weight,  and  red  particles  collect  on  its  surface  ;  but  if  it  is  pro- 
tected from  the  air  by  some  coating  and  then  heated,  there  is  no  gain 
in  weight  and  no  evidence  of  the  red  product.  Therefore,  when  the 
exposed  mercury  is  heated,  something  from  the  air  must  be  added  to  it. 
Now,  if  the  red  substance  is  collected  and  heated  in  a  glass  tube,  mercury 
and  oxygen  are  the  only  products.  Hence,  the  exposed  mercury,  when 
heated,  must  have  combined  with  the  oxygen  of  the  air. 

Oxidation  is  not  always  rapid  enough  to  produce  light 
and  appreciable  heat.  Iron  and  other  metals  rust,  and 
wood  decays  slowly,  but  both  processes  are  mainly  oxida- 
tion. Sometimes  oxidation  develops  considerable  heat. 
Thus,  oily  rags,  piles  of  hay,  and  heaps  of  coal  often  take 
fire  unexpectedly  because  of  the  continued  oxidation.  Such 
oxidation  is  often  called  spontaneous  combustion. 

Substances  which  give  up  oxygen  readily  are  called 
oxidizing  agents.  Potassium  chlorate  is  used  in  fireworks 
for  this  purpose,  and  potassium  nitrate  acts  similarly  in 
gunpowder.  In  the  process  of  oxidation,  oxidizing  agents 


Oxygen.  15 

lose  oxygen,  and  are  said  to  undergo  reduction  —  a  process 
which  will  be  more  fully  described  in  the  next  chapter. 

Oxides  are  formed  when  oxygen  combines  with  other 
elements.  There  are  many  oxides,  and  their  names  express 
in  a  general  way  their  composition.  Oxides  of  different 
elements  are  distinguished  by  placing  the  name  of  the  ele- 
ment (or  a  slight  modification  of  it)  before  the  word  oxide, 
e.g.  magnesium  oxide,  lead  oxide,  zinc  oxide.  Sometimes 
di-,  or  a  similar  numerical  syllable,  is  prefixed  to  the  word 
oxide,  e.g.  carbon  dioxide,  manganese  dioxide,  sulphur 
trioxide,  phosphorus  pentoxide.  The  significance  of  the 
prefix  is  explained  in  Chapter  VII. 

Combustion,  in  a  narrow  sense,  is  rapid  oxidation,  which 
is  always  accompanied  by  light  and  heat.  Popularly,  com- 
bustion means  fire  or  burning,  and  substances  which  burn 
easily  are  called  combustible.  Oxygen  is  essential  to  ordi- 
nary combustion,  and  is  often  called  a  supporter  of  com- 
bustion. Exclude  air  from  a  fire,  and  the  fire  goes  out. 
When  coal  or  wood  burns,  the  carbon  (of  which  they 
largely  consist)  unites  with  the  oxygen  of  the  air,  forming 
thereby  the  invisible  gas  carbon  dioxide,, and  the  chemical 
change  is  manifested  by  heat  and  light.  /Chemically  speak- 
ing, a  substance  burning  in  the  air  is  Uniting  rapidly  with 
oxygen.  But  since  the  air  is  about  one  fifth  oxygen  and 
four  fifths  nitrogen,  —  a  gas  which  does  not  support  com- 
bustion, —  it  follows  that  combustion  is  more  vigorous  in 
oxygen  than  in  air. 

The  correct  explanation  of  fire,  burning,  and  combustion  was  first 
made  by  Lavoisier  (1743-1794).  For  many  years  chemists  had  be- 
lieved that  all  combustible  substances  contained  a  principle  called 
phlogiston,  and  that  when  a  substance  burned,  phlogiston  escaped. 
Very  combustible  substances  were  thought  to  contain  much  phlogiston, 
and  incombustible  substances  no  phlogiston.  This  theory  of  combus- 


1 6  Descriptive  Chemistry. 

tion  was  proposed  by  Becher  (1635-1682)  and  advanced  by  Stahl 
(1660-1734).  Many  famous  chemists —  Priestley,  Scheele,  and  Caven- 
dish—  supported  it.  Lavoisier,  in  1775,  proved  by  his  own  and  others1 
experiments,  that  phlogiston  did  not  exist,  and  that  combustion  is  a 
process  of  combination  with  "  a  certain  substance  contained  in  the  air." 
Soon  after  he  identified  this  substance  as  oxygen.  The  theory  of 
phlogiston,  in  spite  of  its  falsity,  exerted  a  wholesome  influence  on  the 
development  of  chemistry. 

Combustion,  in  a  broad  sense,  is  not  necessarily  oxida- 
tion, but  chemical  action  which  develops  enough  energy 
to  produce  light  and  heat.  This  broader  meaning  will  be 
discussed  later. 

Relation  of  Oxygen  to  Life.  —  Oxygen  is  essential  to  all 
forms  of  animal  and  plant  life.  If  an  animal  or  a  plant  is 
deprived  of  air,  it  dies.  By  respiration  air  is  drawn  into 
the  lungs  and  there  it  gives  up  part  of  its  oxygen  to  the 
blood.  This  oxygen,  which  is  distributed  to  all  parts  of 
the  body  by  trie  blood,  oxidizes  food  and  the  tissues  of  the 
body.  As  a  result  of  this  oxidation  new  tissue  is  built  up 
and  waste  products  are  formed.  One  of  these  waste  prod- 
ucts is  carbon  dioxide  gas,  which  with  other  gases  is 
exhaled  from  the  lungs.  The  blood  during  its  circulation 
turns  dark  red,  owing  to  the  loss  of  oxygen  ;  and  when  this 
dark  red  blood  reaches  the  lungs,  it  receives  a  fresh  supply 
of  oxygen  which  turns  it  bright  red,  thus  preparing  it  for 
another  journey  through  the  body.  Food  must  be  oxidized 
before  it  can  be  taken  up  by  the  body,  and  by  this  oxida- 
tion the  carbonaceous  matter  of  the  body  is  slowly  burned 
to  carbon  dioxide.  It  is  this  slow  oxidation  which  keeps 
the  body  warm.  The  human  body  resembles  a  steam 
engine.  In^each,  the  oxYggiLjQ£-the-air.lielps_burn  fuel  ^ 
largely  composed  of  carbon.  In  the  engine,  the  products 
es"cape  through  a  chimney  and  the  heat  produced  is  used 


Oxygen.  17 

to  form  steam  which  moves  parts  of  the  machine ;  in  the 
body,  the  products  escape  mainly  through  the  lungs  and 
the  heat  keeps  the  body  at  a  temperature  at  which  it  can 
best  perform  its  functions. 

It  was  formerly  believed  that  breathing  pure  oxygen  would  produce 
too  rapid  oxidation  in  the  body  and  burn  up  the  tissue  faster  than  it 
could  be  made.  But  recent  study  shows  that  with  proper  precautions 
oxygen  may  be  breathed  by  a  healthy  person  without  producing  any 
harmful  effect.  The  blood  apparently  absorbs  a  maximum  quantity  of 
oxygen,  whether  supplied  from  air  or  from  the  pure  gas.  Oxygen  is 
often  administered  to  a  person  who  has  been  suffocated,  or  to  one  who 
is  unable  to  inhale  enough  air,  as  in  cases  of  croup,  asthma,  or  extreme 
weakness.  It  is  sometimes  used  to  sustain  life  where  air  is  impure 
or  rare,  as  in  diving  bells  and  submarine  boats,  and  during  balloon 
ascensions  to  a  great  height. 

Decay  is  in  part  oxidation.  The  oxygen  of  the  air 
together  with  water  vapor  acts  upon  animal  and  vegetable 
matter  and  slowly  burns  it  up.  The  decomposition  is  often 
begun  and  hastened  by  bacteria.  The  products  of  decay 
are  numerous,  carbon  dioxide  being  one.  The  oxygen 
dissolved  by  water  assists  in  the  decay  of  the  impurities 
constantly  flowing  into  rivers.  Similarly,  it  oxidizes  in- 
jurious vapors  and  matter  in  the  air,  literally  burning  them 
up,  just  as  it  burns  wood  in  a  stove.  Hence,  running 
water  is  more  likely  to  be  cleaner  than  standing  or  stagnant 
water,  and  the  air  in  the  open  country  or  at  the  seashore 
purer  than  in  the  crowded  city. 

Uses  of  Oxygen.  —  Oxygen  for  commercial  use  is  stored  under 
pressure  in  strong  iron  cylinders.  The  pure  gas  has  limited  use,  since 
air,  although  it  contains  about  80  per  cent  of  the  inert  gas  nitrogen, 
may  usually  be  used  in  place  of  oxygen,  A  mixture  of  oxygen  and 
hydrogen  burned  in  a  suitable  apparatus  produces  an  intensely  hot 
flame,  which  is  sometimes  used  to  melt  refractory  metals  and  to  produce 
the  calcium  light  (see  Oxyhydrogen  Blowpipe). 


1 8  Descriptive  Chemistry. 

Liquid  Oxygen.  —  All  gases  at  a  low  temperature  and 
under  great  pressure  may  be  condensed  to  liquids,  and 
even  to  solids.  Under  these  conditions  oxygen  becomes 
first  a  pale  blue  liquid  and  finally  a  whitish  solid.  A  small 
quantity  was  first  obtained  in  1877,  but  now  it  is  prepared 
by  the  gallon.  It  is  magnetic,  and  when  a  strong  electro- 
magnet is  held  near  its  surface,  the  liquid  suddenly  "leaps 
up  to  the  poles  and  remains  there  permanently  attached 
until  it  evaporates." 

Under  the  normal  pressure  (760  mm.)1  liquid  oxygen  boils  at 
181.4°  C., and  at  this  temperature  its  specific  gravity  is  1. 124  (water  —  i). 

Discovery  of  Oxygen.  —  Oxygen  was  discovered  on 
August  i,  1774,  by  Priestley  (1733-1804).  He  prepared  it 
by  focusing  the  sun's  rays  upon  the  red  mercury  oxide  by 
means  of  "  a  burning  lens  of  twelve  inches'  focal  distance." 
It  was  independently  discovered  by  Scheele  (1742-1786),  a 
Swedish  chemist,  about  the  same  time. 

Priestley  called  the  gas  dephlogisticated  air,  because  he  regarded  it 
as  "  devoid  of  phlogiston."  Scheele  called  it  empyreal  air,  i.e.  fire 
air  or  fire-supporting  air,  because  it  assisted  combustion.  Lavoisier,  in 
1778,  gave  it  the  name  oxygen  (from  the  Greek  oxus,  acid,  and^w,  the 
root  of  a  verb  meaning  to  produce),  because  he  believed  from  his 
experiments  that  oxygen  was  necessary  for  the  production  of  acids  —  a 
view  now  known  to  be  incorrect. 

Weight  of  a  Liter  of  Oxygen.  —  The  volume  occupied 
by  a  gas  depends  upon  the  pressure  and  temperature  to 
which  it  is  subjected.  The  volume  expands  with  rise  of 
temperature  or  with  lowering  of  pressure,  but  contracts 
with  fall  of  temperature  or  with  increase  of  pressure.  In 
general,  if  we  cool  a  gas  or  subject  it  to  a  pressure,  it 
shrinks,  and  if  we  heat  a  gas  or  decrease  the  pressure 

1  This  expression  means  the  normal  or  standard  pressure  of  the  atmosphere  as 
recorded  by  the  barometer  (see  Chapter  VI). 


Oxygen.  19 

it  is  under,  it  expands.  Gas  volumes,  to  be  correctly 
compared,  must  therefore  be  at  the  same  temperature  and 
pressure.  The  normal  or  standard  temperature  is  zero 
degrees  on  the  centigrade  thermometer,  or  briefly  o°  C. 
The  normal  or  standard  pressure  is  the  pressure  of  the 
atmosphere  indicated  by  the  barometer  when  the  mercury  is 
760  millimeters  high,  or  briefly  760  mm.  Under  these 
conditions,  which  are  called  standard  conditions,  a  liter  of 
dry  oxygen  weighs  1.43  gm. 

It  is  not  usually  convenient  to  measure  gases  at  o°  C. 
and  760  mm.  So  if  their  volumes  are  to  be  studied  and 
compared,  it  is  customary  to  reduce  the  observed  volume 
to  the  volume  it  would  occupy  under  standard  conditions. 
This  reduction  is  accomplished  by  applying  two  laws  —  the 
Law  of  Charles  and  the  Law  of  Boyle. 

Law  of  Charles.  —  It  has  been  found  by  experiment  that  under  con- 
stant pressure  all  gases  expand  or  contract  equally  for  equal  changes 
of  temperature.  More  explicitly,  a  gas  expands  or  contracts  ^  of  its 
volume  at  o°  C.  for  every  degree  through  which  it  is  heated  or  cooled. 
This  means  that  273  volumes  at  p°  become  274  at  i°,  275  at  2°,  280  at 
7°,  272  at  —  i°,  270  at  —  3°,  or  273  + 1  volumes  at  /°  (i  e.  at  any  tem- 
perature). This  law  is  not  absolutely  correct,  but  its  variations  from 
the  truth  are  slight. 

Suppose  we  have  10  1.  of  oxygen  at  o°  C.,  and  we  wish  to  know 
the  volume  it  would  occupy  at  15°  C.  The  problem  is  easily  solved  by 
stating  it  as  a  proportion,  thus  — 

273:273+  I5::lo:.r.    - 

The  value  of  .ris  the  volume  required.  Conversely,  in  reducing  10  vol- 
umes at  15°  C.  to  the  volume  occupied  at  o°  C.,  the  proportion  is  — 

273  +  I5:273::io:;r. 

If  the  given  temperature  is  below  o°,  the  number  of  degrees  is  subtracted 
from  273. 

Law  of  Boyle.  —  It  has  also  been  found  by  experiment  that  under 
constant  temperature  the  volume  of  a  gas  is  inversely  proportional  to 


lo  Descriptive  Chemistry. 

the  pressure.  This  is  Boyle's  law.  It  means  that  doubling  the  pres- 
sure halves  the  volume,  and  vice  "versa.  Like  the  above  law,  this  law  is 
only  approximately  correct. 

Suppose  we  have  10  1.  of  oxygen  at  760  mm.,  and  we  wish  to  know 
the  volume  it  would  occupy  at  775  mm.  According  to  the  law,  the 
proportion  expressing  the  relation  is  — 

760:775::^:  10. 

The  value  of  x\<&  the  required  volume.  Conversely,  if  we  have  10  1.  at 
775  mm.,  and  wish  to  know  its  standard  volume,  the  proportion  is  — 


It  is  convenient  to  notice  that  the  proportion  is  stated  so  that  the 
extremes  (or  means)  are  the  original  pressure  and  volume.  In  other 
words,  one  pressure  multiplied  by  its  volume  equals  the  other  pressure 
multiplied  by  its  volume,  or  — 

P\P\\V\  V. 

Hence,  the  proportion  is  applicable  to  values  not  necessarily  includ- 
ing 760. 

EXERCISES. 

i  .   What  is  the  symbol  of  oxygen  ? 

2.  How  is  oxygen  prepared  (a)  in  the  laboratory,  and  (£)  commer- 
cially ? 

3.  Name  several  compounds  from  which  oxygen  can  be  prepared. 

4.  Summarize  the  properties  of  oxygen.     What  is  its  most  charac- 
teristic property  ? 

5.  If  air  contains  something  besides  oxygen,  what  must  be  the  gen- 
eral properties  of  this  other  ingredient  ? 

6.  Define  and  illustrate  (a)  oxidation,  (ft)  oxide,  (c)  combustion, 
(//)  oxidizing  agent. 

7.  What   elements   were   mentioned   in    studying   oxygen  ?     What 
compounds  ? 

8.  What  general  chemical  change  is  involved   in  burning  ?     What 
class  of  chemical  changes  is  illustrated  by  (a}  preparation  of  oxygen 
from  mercuric  oxide,  (b}  burning  of  sulphur  in  oxygen  ? 

9.  Give  a  brief  account  of  Priestley,   Scheele,  and  Lavoisier  (see 
Appendix,  §  4). 


Oxygen.  21 

10.  What  chemical  part  does  oxygen 'take  in  (a)  respiration,  (#)  de- 
cay, (c)  combustion,  («)  oxidation  ? 

11.  State  and  illustrate  (a)  Charles's  law  and  (£)  Boyle's  law. 

12.  Give  a  brief  account  of  Boyle  and  of  Charles. 


PROBLEMS. 

1.  Potassium  chlorate  contains  about  39  per  cent  of  oxygen.     How 
many  grams  of  oxygen  can  be  prepared  from  (rt)  100  gm.,  (^)  250  gm., 
and  (c)  725  gm.  of  potassium  chlorate  ? 

2.  What  approximate  weight  of  oxygen  can  be  prepared  from  100 
gm.  of  potassium  chlorate  containing  12  per  cent  of  impurity  ? 

3.  What  is  the  weight  of  (a)   10  1.  of  oxygen,  (b)  75-!.,  (c)  500 
cc.,  (d)  750  cc.,  0)  4!.? 

4.  A  room  25  m.  long,  17  wide,  and  15  high  is  filled  with  oxygen. 
What   weight    of  gas   does   it   contain  ?    (A   liter   of  oxygen   weighs 

1-43  gm.) 

5.  Reduce  the  following  volumes  to  the  volume  occupied  at  o°  C. : 
(a)  173  cc.  at  12°  C.,  (b)  466  cc.  at  14°  C.,  (c)  706  cc.  at  15°  C.,  (d) 
25  cc.  at  27°  C. 

6.  A  volume  of  gas  at  o°  C.  measures  1500  cc.    What  is  its  volume 
at  (a}  15° C.,  (d)  50°  C.,  (0  100° C.,  (d}  300° C.  ? 

7.  If  500  cc.  of  gas  at  27°  C.  are  cooled  to  —  5°C.,  what  is  the  new 
volume  ? 

8.  Reduce    the    following    volumes   to   the   volume    occupied   at 
760  mm.  :     (a)  200  cc.  at  740  mm.,  (b)  25  cc.  at  780  mm.,  (c)  467  cc. 
at  756  mm.  Ans.  (a)  1947?  (^)  25.65,  (c)  464-54- 

9.  A  gas  measures  1000  cc.  at  770  mm.     What  is  its  volume  at 
530  mm.? 

10.  Reduce  the  following  to  standard  conditions:  (a)  147  cc.  at 
570  mm.  and  136.5°  C.,  (b}  320  cc.  at  950  mm.  and  9i°C,  (c)  480  cc. 
at  380  mm.  and  68.25°C,  (d)  25  cc.  at  780  mm.  and  27°  C.,  (*)  14  cc. 
at  763  mm.  and  ii°C. 

Ozone  is  a  gas  related  to  oxygen,  though  its  properties  differ.  It  is 
formed  when  electric  sparks  pass  through  the  air,  and  is  therefore  pro- 
duced when  electrical  machines  are  in  operation  and  during  thunder 
storms.  Slow  oxidation,  especially  of  moist  phosphorus,  produces 
ozone.  Indeed,  its  formation  accompanies  several  chemical  changes. 


22  Descriptive  Chemistry. 

such  as  the  burning  of  hydrogen  and  of  certain  resins,  and  the  decom- 
position of  water  by  electricity. 

Ozone  has  a  peculiar  odor,  suggesting  burning  sulphur.  The  name 
ozone  signifies  smell.  It  is  active  chemically,  tarnishing  metals,  bleach- 
ing colored  vegetable  substances,  deodorizing  foul  animal  matter,  and 
corroding  such  substances  as  cork  and  rubber.  It  is  sometimes  used  as  a 
disinfectant,  though  other  oxidizing  agents  are  more  convenient.  When 
heated  to  250°  C.,  or  higher,  it  is  wholly  changed  into  oxygen.  Ozone, 
therefore,  contains  nothing  but  oxygen.  When  oxygen  is  changed  into 
ozone,  it  is  found  that  three  volumes  of  oxygen  yield  two  volumes  of 
ozone  ;  and,  conversely,  the  two  volumes  of  ozone,  when  heated,  become 
three  volumes  of  oxygen.  Hence,  volume  for  volume,  ozone  is  1.5  times 
heavier  than  oxygen.  For  this  reason  ozone  is  sometimes  called  "con- 
centrated oxygen,"  or  "an  oxide  of  oxygen."  Its  theoretical  relation  to 
oxygen  will  be  subsequently  discussed. 

The  atmosphere  usually  contains  a  small  proportion  of  ozone,  prob- 
ably not  more  than  one  volume  in  700,000  volumes  of  air.  It  is  more 
abundant  in  the  open  country  and  at  the  seashore  than  in  cities. 


CHAPTER   III. 
HYDROGEN. 

Occurrence.  —  Free   hydrogen  is  present  in  the  gases 

petroleum  wells,  and  natural 


gas  openings.     Artificial  illuminating  gas  contains  consid-  ^ 
erable  hydrogen.      It  is^T  product  of  fermentation  and  *»"* 
decay,  and  according  to  recent  observations  a  very  small 
quantity  is  present  in  the  atmosphere  of  the  earth.     Enor-  *— 
mous  quantities  of  free  hydrogen  exist  in  tne  atmqsrjhere  7 
of  the   sun,  and   during   an   eclipse  of   the   sun  gigantic 
streams  of  burning  hydrogen  may  be  seen  shooting  out 
from  the  sun's  disk  thousands  of  miles  into  space.     Other 
heavenly  bodies  which  are  self-luminous,  like  the  star  Sirius 
and  the  nebulae,  contain  free  hydrogen.     The  spectroscope 
has  revealed  its  presence  in  these  distant  bodies.     Meteor-*? 
ites^  which  come  from  regions  far  beyond  our  earth,  often 
contain  free  hydrogen. 

Cojnbinedji^drogen  is  abundant  and  widely  distributed.  * 
It  forms  one  ninthly  weight^  of  water.     Most  animal  and  f  ^ 
vegetable  matter  contains  hydrogen.     It  is  also  an  essential  1  1 
component  nf_a1J_gHHs      Combined  with  carbon,  it  forms 
many  gases  and  liquids  called  hydrocarbons,  which  are  con-  '   * 
stituents  of  illuminating  gas,  kerosene,  and  naphtha.    Com- 
bined with  carbon  and  Oxygen,  It  forms  many  vegetable  >  C> 
compounds,  such  as  sugar,  starch,  parser,  wood,  and  numer- 
ous artificia'1  products.    With  nitrogen  it  forms  the  familiar  ^ 
compound,  ammonia  ;  and  with  sulphur,  the  bad-smelling  gas,     ^ 
hydrogen  sufphTdeTwhich  occurs  in  many  sulphur  springs.  ' 

23 


24  Descriptive  Chemistry. 

Preparation.  —  Hydrogen,  like  oxygen,  is  prepared  from 
its  compounds.  In  the  laboratory  this  is  easily  accom- 
plished  by  allowing  a  metal  and  an  acid  to  interact.  The 
metals  usually  employed  are  zinc,  iron,  or  magnesium,  and 
the  acids  are  dilute  sulphuric  acid  or  hydrochloric  acid. 
The  hydrogen  comes  from  the  acid  and  bubbles  through 
the  liquid,  when  the  acid  and  metal  are  put  into  a  test  tube 
or  flask.  On  a  large  scale  hydrogen  is  prepared  in  a  genera- 
tor, which  consists  of  a  glass  vessel  provided  with  a  delivery 
tube  arranged  to  collect  the  gas  over  water  in  a  pneumatic 
trough.  No  flame  should  be  near  during  the  performance 
of  this  experiment,  because  mixtures  of  air  and  hydrogen 
explode  violently  when  ignited.  The  interaction  of  zinc 
and  sulphuric  acid  produces,  besides  hydrogen,  a  compound 
called  zinc  sulphate.  This  remains  in  the  generator  in 
solution,  and  if  the  solution  is  allowed  to  evaperate,  the 
zinc  sulphate  separates  as  transparent  crystals,  which  soon 
turn  white  in  the  air.  Hydrogen  may  be  obtained  from 
water  by  allowing  tiie_Jii£lal-ao^lijLir^^  to  interact. 


If  a  small  piece  of  sodium  is  dropped  upon  cold  water,  the  sodium 
melts  into  a  shining  globule,  which  spins  about  rapidly  on  the  water 
with  a  hissing  sound,  and  finally  disappears  with  a  slight  explosion. 
But  when  the  sodium  is  wrapped  in  a  piece  of  tea  lead  pierced  with  a 
few  holes  and  then  dropped  beneath  the  shelf  of  a  pneumatic  trough 
filled  with  water,  the  action  proceeds  smoothly.  Hydrogen  gas  rises 
and  displaces  the  water  from  a  test  tube  or  bottle  supported  over  the 
hole  in  the  shelf.  The  nature  of  the  chemical  change  which  attends 
the  liberation  of  hydrogen  from  water  will  be  explained  later  (Chap- 
ter V). 

Hydrogen,  together  with  oxygen,  is  liberated  from  water 
by  passing  a  current  of  electricity  througlTwafer  containing 
a  little  sulphuric  acid  (see  Chapter  V). 

Hydrogen  may  also  be  prepared  by  passing  steam  —  the 
gaseous  form  of  water  —  over  heated  metals. 


Hydrogen.  25 

This  experiment  was  first  performed  by  Lavoisier,  in  1783,  while  he 
was  studying  the  composition  of  water.  He -passed  steam  through  a 
red-hot  gun  barrel  containing  bits  of  iron.  The  oxygen  of  the  steam 
combined  with  the  iron,  and  the  hydrogen  escaped  from  the  tube.  Since 
Lavoisier  was  studying  the  composition  of  water,  and  not  the  properties 
of  hydrogen,  he  naturally  thought  of  this  gas  as  essential  for  forming 
water.  So  he  says  in  his  notes,  "  No  name  appears  to  us  more  suitable 
than  that  of  hydrogen,  that  is  to  say,  'generative  principle  of  water.'" 
Apart  from  historical  interest,  this  experiment  has  commercial  value. 
If  steam  is  passed  over  red-hot  coal  (instead  of  iron),  producer  gas 
is  formed.  This  is  a  mixture  consisting  largely  of  hydrogen,  which  is 
used  as  a  source  of  heat  in  making  steel  and  glass.  If  oil  vapor  is 
added  to  this  mixture,  water  gas  is  formed.  This  is  an  illuminating 
gas  like  ordinary  illuminating  gas,  and  is  used  in  many  cities  (see 
Water  Gas). 

Physical  Properties.  —  Hydrogen  has  no  taste  or  color. 
The  pure  gas  has  no  odor,  though  hydrogen  as  ordinarily 
prepared  has  a  disagreeable  odor,  due  mainly  to  impurities 
in  the  metals  used.  Most  of  these  impurities  may  be  re- 
moved by  passing  the  gas  through  a  solution  of  potassium 
permanganate.  Hydrogen  is-the  lightest  known  substance. 
One  liter  of  dry  hydrogen  at  o°  C.  and  760  mm.  weighs 
only  0.0896  gm.  Volume  for  volume,  air  is  about  14.4 
times,  oxygen  16  times,  and  water  11,000  times  heavier 
than  hydrogen. 

The  extreme  lightness  of  hydrogen  may  be  easily  shown,  (i)  If  a 
wide-mouth  bottle  of  the  gas 
is  left  uncovered  two  or  three 
minutes  and  a  lighted  match 
then  dropped  in,  the  match 
will  continue  to  burn.  If 
hydrogen  had  been  present, 
the  flame  would  have  caused 
it  to  combine  with  the  oxy- 

gen  of  the  air  with   a  loud  FlG   lelpouring  hydrogen. 

explosion.    (2)  If  a  bottle  of 
hydrogen  is  held  beneath  a  bottle  of  air  as  shown  in  Figure  i,  the  gases 


26  Descriptive  Chemistry. 

soon  exchange  places,  the  hydrogen,  owing  to  its  lightness,  rising  into 
the  upper  bottle.  Its  presence  there  may  be  readily  shown  by  dropping 
a  lighted  match  into  this  bottle ;  if  the  experiment  has  been  well  done, 
the  hydrogen  will  burn,  but  in  most  cases  the  loud  explosion  shows 
that  only  a  part  of  the  hydrogen  has  been  poured  upward.  A  lighted 
match  dropped  into  the  other  bottle  reveals  only  air.  (3)  If  a  small 
collodion,  or  rubber,  balloon  is  filled  with  hydrogen  and  then  released, 
it  will  rise  rapidly  into  the  air.  Hydrogen,  because  of  its  lightness,  is 
sometimes  used  to  fill  large  balloons,  but  ordinary  illuminating  gas  is 
usually  employed. 

Hydrogen  is  the  standard  for  reckoning  the  density  of  gases.  Thus, 
since  a  liter  of  oxygen  weighs  1.43  gm.,  its  density  is  found  by  the 
proportion:-  Q  ^ .  ,  ^  . .  ,  .  ^  .  ^  l6 

Hydrogen  is  not  very  soluble  in  water,  but  it  is  absorbed 
by  several  metals,  especially  the  rare  metal  palladium. 
This  property  of  absorbing  gases  is  called  occlusion. 

Only  about  1.84  1.  of  hydrogen  at  760  mm.  pressure  dissolve  in 
100  1.  of  water  at  20°  C.  Palladium  absorbs  from  370  to  960  times 
its  own  volume  of  hydrogen,  according  to  the  conditions  of  the  experi- 
ment. Platinum  and  iron  act  similarly,  though  to  a  less  degree.  Illu- 
minating gas,  which  contains  considerable  hydrogen,  is  also  absorbed 
by  metals.  And  since  heat  is  developed  by  occlusion,  the  illuminating 
gas  may  be  lighted  by  the  heated  metal  upon  which  it  flows.  A  self- 
lighting  gas  burner  acts  on  this  principle.  The  act  of  occlusion  is 
partly  chemical  and  partly  physical. 

Hydrogen  illustrates  diffusion;  i.e.  it  readily  passes 
through  porous  substances  and  completely  mixes  with 
other  gases  without  stirring  or  agitating. 

It  penetrates  unglazed  earthenware,  paper,  and  heated  metals,  espe- 
cially platinum.  Hydrogen  has  the  highest  rate  of  diffusion,  because 
its  density  is  the  lowest.  The  rate  of  diffusion  of  a  gas  is  inversely 
proportional  to  the  square  root  of  the  density.  Thus,  the  rate  of  diffu- 
sion of  hydrogen  is  four  times  that  of  oxygen,  since  the  density  of  oxy- 
gen is  sixteen  times  that  of  hydrogen.  We  are  largely  indebted  for 
our  knowledge  of  diffusion  to  the  English  chemist,  Thomas  Graham 
(1805-1869). 


Hydrogen.  27 

Hydrogen  is  not  poisonous  if  pure.  It  does  not  sup- 
port life,  but  a  little  may  be  breathed  without  danger. 
When  the  lungs  are  filled  with  it  the  voice  becomes  very 
shrill  and  thin. 

Chemical  Conduct.  —  Hydrogen  burns  in  the  air  and 
in  oxygen  with  an  almost  invisible  but  very  hot  flame. 
Water  is  the  product  of  its 
combustion.  These  facts  may 
be  verified  by  the  apparatus 
shown  in  Figure  2.  The  hydro- 
gen, which  is  generated  from 
zinc  and  hydrochloric  acid  in 
the  flask,  passes  through  the 
U-tube  filled  with  calcium 
chloride  (to  remove  the  mois- 
ture), and  is  lighted  at  the  tip 
after  it  has  driven  all  the  air  from  the  apparatus.1  A 
platinum  or  copper  wire  held  in  the  flame  instantly  becomes 
red-hot.  If  a  small,  dry,  cold  bottle  is  held  over  the  flame, 
moisture  is  deposited  inside  the  bottle. 

The  film  of  water  often  noticed  on  the  bottom  of  a  vessel  placed 
over  a  lighted  gas  range  or  a  Bunsen  burner  is  formed  by  the  burning 
hydrogen  and  hydrogen  compounds  of  the  illuminating  gas.  Similarly, 
water  often  drops  from  the  top  of  the  oven  of  a  lighted  gas  range.  Or- 
ganic substances  containing  hydrogen,  such  as  wood  and  paper,  when 
burned,  yield  water  as  one  of  their  products. 

The  fact  that  the  only  product  of  burning  hydrogen  is  water  was  first 
shown  in  1783  by  Cavendish  (1730-1810).  Lavoisier  in  the  same  year 
verified  this  fact  and  utilized  it  to  explain  the  composition  of  water. 

The  temperature  of  the  hydrogen  flame  is  very  high. 
More  heat  is  produced  by  burning  hydrogen  in  oxygen 


FIG.  2.  —  Apparatus  for  burning 
hydrogen.   . 


1  This   experiment   is   dangerous.     The   precautions   to   be   observed   can   be 
found  on  pages  48-49  in  the  author's  "  Experimental  Chemistry." 


28  Descriptive  Chemistry. 

than  by  burning  the  same  weight  of  any  other  substance 
(see  Chapter  X). 

Hydrogen  burns  in  chlorine  gas.  The  flame  is  bluish  white,  not 
very  hot,  and  the  product  is  hydrochloric  acid  gas  —  a  compound  of 
hydrogen  and  chlorine.  This  burning  of  hydrogen  in  chlorine  illus- 
trates the  broader  use  of  the  word  combustion,  since  no  oxygen  is 
involved. 

Hydrogen  does  not  support  combustion,  as  the  term  is 
usually  used.  This  fact  is  illustrated  by  putting  a  lighted 
taper  into  an  inverted  bottle  of  hydrogen.  The  taper 
ignites  the  hydrogen,  which  burns  at  the  mouth  of  the 
bottle.  The  taper  does  not  burn  inside  the  bottle,  but  when 
it  is  slowly  withdrawn  through  the  burning  hydrogen  it  is 
relighted.  Hence,  hydrogen  burns,  but  does  not  support 
combustion. 

A  mixture  of  hydrogen  and  air  explodes  violently  when 
ignited.  Therefore,  the  air  should  be  fully  expelled  from 
the  apparatus  in  which  hydrogen  is  being  generated  before 
the  gas  is  collected,  and  no  flames,  large  or  small,  should  be 
near.  Neglect  of  these  precautions  has  caused  serious 
accidents. 

Hydrogen  not  only  combines  energetically  with  frea 
oxygen,  but  it  withdraws  oxygen  from  compounds.  As 
stated  before,  this  chemical  removal  of  oxygen  is  called 
reduction.  Hydrogen  is  a  vigorous  reducing  agent. 

The  Oxyhydrogen  Blowpipe  utilizes  the  intense  heat  pro- 
duced by  burning  a  mixture  of  hydrogen  and  oxygen.  The 

apparatus  (Fig.  3)  con- 
sists of  two  pointed  metal 
tubes.  The  inner  and 
smaller  one  is  for  the 
Blowpipe  tip.  oxygen,  and  the  outer 

and  larger  one  for  the  hydrogen.     Their  pointed  ends  are 


Hydrogen.  29 

close  together,  and  the  two  gases  mix  as  they  are  forced 
out  of  these  small  openings  by  the  pressure  maintained  in 
the  storage  tanks.  Sometimes  the  tubes  are  separated, 
but  the  gases  flow  from  a  similar  opening.  The  hydrogen 
is  first  turned  on  and  lighted  at  the  pointed  opening ;  then 
the  oxygen  is  turned  on  and  the  flow  gradually  regulated 
until  the  flame  is  the  desired  size,  usually  thin,  straight, 
and  as  long  as  the  apparatus  requires.  There  is  no  danger 
in  using  the  blowpipe,  provided  it  does  not  leak  and  the 
pressure  is  properly  regulated  by  the  stopcocks.  In  the 
hot  flame,  some  metals,  like  silver,  turn  to  vapor ;  some, 
like  iron,  burn  brilliantly ;  while  others,  like  platinum,  melt. 
When  the  flame  strikes  against  a  piece  of  lime  of  other  sub- 
stance difficult  to  melt,  the  lime  becomes  intensely  bright. 
Thus  used,  it  is  called  the  lime,  calcium,  or  Drummond 
light  and  is  often  employed  in  operating  the  stereopticon. 

The  blast  lamp  is  a  modification  of  the  oxyhydrogen  blowpipe.  The 
apparatus  (Fig.  4)  consists  of  two  tubes,  an  inner  one  for  air  and  an 
outer  one  for  illuminating  gas.  The  air, 
which  is  forced  through  the  apparatus  by 
a  bellows,  provides  oxygen,  and  the  illumi- 
nating gas  contains  hydrogen  and  other 
combustible  gases.  The  mixture  burns  at 
the  opening  of  the  tubes  with  a  colorless 
or  bluish  flame,  which  is  hotter  than  the 
Bunsen  flame  —  the  usual  source  of  heat  for 
chemical  experiments.  The  shape  of  the 
flame  is  easily  regulated  by  stopcocks. 

Liquid  Hydrogen  is  a  colorless,  trans- 
parent liquid  produced  bv  subjecting  the 

FIG.  4.  — Blast  lamp, 
gas  to  great  pressure  and  low  temperature. 

It  was  first  produced  in  1898  by  Dewar.  The  temperature  used  was 
—  205°  C.,  and  the  pressure  was  180  atmospheres  (i.e.  180  times  760 
mm.).  At  the  ordinary  pressure  it  boils  at  —  238°  C.  Under  reduced 
pressure  and  at  —  256°  C.  it  becomes  "a  white  mass  of  solidified  foam." 


jo  Descriptive  Chemistry. 

Discovery  of  Hydrogen.  —  Paracelsus  in  the  sixteenth  century  ob- 
tained hydrogen  by  the  interaction  of  acids  and  metals.  It  was  iden- 
tified as  an  element  in  1766  by  Cavendish,  who  called  it  inflammable 
air.  The  name  hydrogen,  given  to  it  by  Lavoisier,  in  1783,  is  derived 
from  the  Greek  words  hudor,  water,  and  gen,  the  root  of  a  verb  mean- 
ing to  produce. 

EXERCISES. 

1.  What  is  the  symbol  of  hydrogen  ? 

2.  What  familiar  compounds  contain  hydrogen? 

3.  How  is  hydrogen  prepared  in  the  laboratory?     Describe  other 
methods  of  preparation. 

4.  Summarize  the  properties  of  hydrogen.     What  is  its  most  char- 
acteristic property  ? 

5.  Why  is  there  danger  of  an  explosion  in  generating  hydrogen? 
How  may  the  danger  be  avoided  ? 

6.  What  is  the  weight  of  a  liter  of  dry  hydrogen?      How  many 
times  heavier  than  a  liter  of  hydrogen  is  one  of  air  ? 

7.  Define  and  illustrate  (a)  occlusion  and  (b}  diffusion  of  gases. 

8.  What  chemical  change  occurs  when  hydrogen  burns  in  air  ? 

9.  Is  water  an  oxide  ?     Why  ? 

10.  How  does  the  heat  of  the  hydrogen  flame  compare  with  its 
luminosity  ? 

n.  Define  (#)  reduction  and  (£)  reducing  agent.  Name  a  reduc- 
ing agent. 

12.  Describe  («)  the  compound   blowpipe  and  (&)  the   blast  lamp, 
and  state  the  use  of  each. 

13.  Summarize  briefly  the   discovery  of  hydrogen.      Give  a  short 
account  of  Cavendish.    Why  and  by  whom  was  hydrogen  so  named  ? 

14.  What  class  of  chemical  changes  is  illustrated  by  («)  the  prepara- 
tion  of  hydrogen  from  zinc  and  sulphuric  acid,  (<£)  the  burning   of 
hydrogen  in  air  ? 

PROBLEMS. 

1.  How  many  times  heavier  than  a  liter  of  hydrogen  is  a  liter  of 
oxygen,  both  being  dry  and  under  standard  conditions  ? 

2.  What  is  the  weight  of  (a)  500  cc.  of  dry  hydrogen  gas  at  o°  C. 
and  760  mm.  ?     (b)  Of  1800  cc.  ?     (V)  Of  9  1.  ? 

3.  The  standard  pressure  at  which  a  gas  is  measured  is  760  mm. 
Express  the  same  in  inches. 


CHAPTER  IV. 
GENERAL  PROPERTIES  OF  WATER. 

WATER  is  worthy  of  extensive  study  because  of  its 
importance  in  the  animal,  vegetable,  and  mineral  king- 
doms, its  peculiar  properties,  and  its  numberless  uses. 

Occurrence  in  Nature. — Water,  in  the  form  of  vapor, 
is  always  present  in  the  atmosphere.  Evaporation  is  con- 
stantly taking  place  from  the  surface  of  the  ocean,  from 
the  moist  earth,  from  the  bodies  of  animals,  and  from 
plants.  This  vapor  is  continually  condensing,  and  appears 
as  clouds,  mist,  fog,  rain,  snow,  hail,  dew,  and  frost. 

The  proportion  of  water  vapor  in  the  atmosphere  varies  between  wide 
limits,  the  amount  present  being  largely  influenced  by  the  temperature. 
It  has  been  found,  however,  that  1000  volumes  of  ordinary  air  contain 
about  14  volumes  of  water  vapor.  The  total  amount  of  vapor  in  the  atmos- 
phere is  beyond  comprehension. 

In  the  liquid  state  water  occurs  in  vast  quantities. 
About  three  fourths  of  the  surface  of  the  globe  is  covered 
with  water.  Soil  and  porous  rocks  hold  considerable 
quantities,  and  plants  and  animals  contain  a  large  pro- 
portion. Many  substances  which  are  apparently  dry  really 
contain  a  large  proportion  of  water.  Thus,  in  a  ton  of 
clover  hay  there  are  upwards  of  200  Ib.  of  water,  and  a 
ton  of  salt  hay,  which  is  usually  very  dry,  contains  about 
100  Ib. 

Many  common  foods  are  largely  water,  as  may  be  seen 
by  the  following  — 

3' 


Descriptive  Chemistry. 


TABLE  OF  THE  PROPORTION  OF  WATER  IN  FOOD. 


FOOD. 

PER  CENT 
OF  WATER. 

FOOD. 

PER  CENT 
OF  WATER. 

Cod                    .... 

8^.6 

Q4..3 

Beef                            .     . 

6l.Q 

Apples  

84.6 

Lobster 

"•'y 

7Q  2 

Strawberries        .     . 

QO  A. 

Ecrorg    . 

/y  ••* 
T\-1 

Watermelon  .... 

yw"4- 
02.4. 

Asparagus 

04.. 

Milk  

87. 

Potatoes 

78.7 

Cheese  ....... 

28  to  72 

Cucumbers  

954 

White  bread  .... 

35-3 

The  human  body  is  nearly  70  per  cent  water,  and  during  a 
year  the  average  man  drinks  about  half  a  ton. 

Water  in  the  form  of  ice  permanently  covers  the  coldest 
parts  of  the  surface  of  the  earth,  e.g.  the  polar  regions  and 
the  summits  of  high  mountains.  A  rough  estimate  of  the 
total  weight  of  ice  on  the  earth's  surface  is  6,373,000,0x30 
millions  of  metric  tons.1 

Functions  of  Water  in  Nature.  —  Since  water  is  the 
only  liquid  occurring  in  large  quantities  on  the  earth's  sur- 
face, it  is  the  great  agent  of  erosion.  It  cuts  away  the 
earth's  crust,  and  transports  the  material  from  higher  to 
lower  levels,  or  washes  it  into  the  ocean.  Together  with 
carbon  dioxide  gas  it  decomposes  the  rocks,  changing  them 
into  clay,  sand,  and  substances  which  make  the  soil  pro- 
ductive. Its  cycle  of  changes  from  liquid  to  vapor  and 
vapor  to  liquid  exerts  a  marked  influence  on  the  distribu- 
tion of  heat  and  moisture  upon  the  earth's  surface,  i.e.  on 
climate. 

It  dissolves  many  solids  and  gases  and  is  constantly  re- 
moving from  the  rocks  and  soil  their  soluble  constituents, 


1  A  metric  ton  contains  2204.6  pounds. 


Properties  of  Water.  33 

some  of  which  serve  for  the  nutrition  of  plants,  though  the 
larger  part  passes  on  to  the  ocean.  The  latter  thus  be- 
comes a  vast  reservoir  of  water  containing  salt  and  other 
mineral  matter  obtained  from  the  earth's  crust.  In  the 
vital  processes  of  animals  and  plants  it  helps  change  the 
food  into  a  condition  fit  for  distribution  and  assimilation. 

Industrial  Applications.  —  Besides  the  universal  use  of 
water  for  drinking,  it  is  applied  to  an  endless  variety  of  use- 
ful and  convenient  purposes.  It  has  always  been  man's 
beast  of  burden.  It  is  the  vehicle  for  transferring  mechan- 
ical energy  to  water  wheels  —  an  application  now  being 
made  on  a  vast  scale  for  generating  electricity.  It  utilizes 
by  its  peculiar  properties  the  energy  in  fuel  by  means  of 
the  steam  engine.  It  is  the  highway  for  transportation  on 
the  largest  scale  by  ocean,  river,  lake,  and  canal.  It  is  the 
vehicle  for  the  distribution  of  heat  by  hot  water  and  steam. 
It  is  the  indispensable  solvent  in  metallurgy,  in  the  manu- 
facture of  chemicals,  and  in  such  industries  as  soap 
making,  bleaching,  brewing,  dyeing,  and  tanning;  it  is 
necessary  wherever  mortar  and  cement  are  used.  Man's 
work  would  be  stopped  in  a  thousand  other  ways  were 
he  deprived  of  water. 

Physical  Properties  of  Pure  Water.  —  Owing  to  its 
remarkable  solvent  power,  water  is  never  found  pure  in 
nature,  and  is  purified  even  in  the  laboratory  only  by  taking 
especial  precautions.  At  the  ordinary  temperature  water 
is  a  tasteless  and  odorless  liquid.  It  is  usually  colorless, 
but  thick  layers  are  bluish.  Water  is  a  poor  conductor  of 
heat. 

This  last  property  may  be  shown  by  boiling  water  near  the  surface 
in  a  large  test  tube  containing  a  piece  of  ice  weighted  down  upon  the 
bottom.  The  ice  remains  unmelted  for  some  time,  although  the  water 
is  boiling  a  few  inches  above  it. 


34  Descriptive  Chemistry. 

Most  liquids  expand  with  heat  and  contract  with  cold. 
Water  is  an  exception.  If  water  at  100°  C.  is  gradually 
cooled,  it  contracts  in  volume.  But  when  4°  C.  is  reached, 
if  the  cooling  continues,  the  volume  increases  as  long  as 
the  liquid  state  is  maintained.  Hence  at  4°  C.  a  given 
volume  contains  the  greatest  weight  of  water.  That  is, 
water  has  its  maximum  density  at  4°  C. 

The  density  of  water  at  4°  C.  is  i ;  and  water  at  this  temperature  is 
the  standard  for  determining  the  densities  of  solids  and  liquids.  Thus, 
when  we  say  the  density  of  gold  is  19,  we  mean  that  gold  is  19  times 
heavier  than  an  equal  volume  of  water  at  4°C. 

The  expansion  of  water  when  cooled  from  4°  C.  to  o°  C.  is  slight,  but 
the  change  is  exceedingly  important  in  nature.  When  the  water  on  the 
surface  of  a  lake  or  river  cools,  it  contracts,  and  since  it  is  heavier 
(volume  for  volume)  than  the  warmer  water  beneath,  it  sinks.  The 
warmer  water  rises,  is  cooled,  and  likewise  sinks,  thus  causing  a  circula- 
tion which  continues  until  all  the  water  from  surface  to  bottom  has  the 
temperature  of  4°C.  Now  if  the  cooling  continues,  the  surface  water 
expands  and  remains  on  top,  because  it  is  lighter  than  the  water 
beneath.  Hence  when  the  temperature  of  the  air  falls  to  o°C,  this  top 
layer  of  water  freezes  and  protects  the  remaining  water  from  the  cold, 
thus  stopping  the  circulation.  Should  the  circulation  continue,  as  the 
temperature  fell  from  4°  C.  to  o°  C.,  the  whole  body  of  water  would 
finally  freeze  from  top  to  bottom.  This  condition  would  not  only 
destroy  the  fish  and  marine  plants,  but  seriously  affect  climate,  since 
the  heat  of  summer  could  not  melt  such  a  vast  mass  of  ice. 

When  water  freezes,  it  expands  about  one  tenth  of  its 
volume.  That  is,  100  cc.  of  water  produce  about  no  cc. 
of  ice.  In  other  words,  100  cc.  of  water  and  1 10  cc.  of  ice 
weigh  100  gm.  Hence  ice  floats.  The  specific  gravity  of 
ice  is  about  0.92. 

The  pressure  exerted  by  water  when  it  freezes  is  powerful.  Vessels 
or  pipes  completely  filled  with  water  often  burst  when  the  water  freezes. 
It  is  an  erroneous  but  popular  idea  that  "  thawing  out "  a  pipe  bursts  it. 
As  a  matter  of  fact,  ice  contracts  when  it  melts.  The  pipe  cracks  when 
the  water  freezes,  and  as  the  ice  melts  a  channel  is  left  for  the  water  to 


Properties  of  Water.  35 

flow  out  of  the  pipe.  Because  of  this  property,  ice  is  an  effective  agent 
in  splitting  rocks.  Water  creeps  into  cracks,  especially  into  the  narrow 
ones  by  capillary  attraction,  and  when  it  freezes,  the  rock  is  slowly  split 
apart.  Water  in  freezing  also  destroys  the  tissue  of  living  plants,  which 
are  often  said  to  have  been  "touched  by  frost."  Frozen  flesh  for  a 
similar  reason  becomes  pulpy  and  is  more  liable  to  putrefy  when  thawed 
—  a  fact  sometimes  overlooked  by  those  who  eat  flesh  food  which  has 
been  kept  in  cold  storage. 


FIG.  5.  —  Snow  crystals. 
From  photographs  by  Wilson  A.  Bentley. 


Ice  melts  at  o°  C.  (32°  F.),  which  is  also  the  freezing 
point  of  water.     Ice  often  crystallizes  in  freezing,  but  the 


36  Descriptive  Chemistry. 

individual  crystals  are  seldom  visible  except  during  the  first 
stages  of  the  process.  Snow  crystals  are  common  (Fig.  5). 
They  are  always  six-sided,  and  are  formed  in  the  atmos- 
phere by  the  freezing  of  water  vapor. 

Water  evaporates  at  all  temperatures,  passing  off  as  an 
invisible  vapor  into  the  atmosphere  or  into  the  air  confined 
over  it.  If  water  is  heated,  the  vapor  passes  off  rapidly 
until  the  thermometer  reads  ioo°C.  (or  212°  F.).  At  this 
point  water  boils,  i.e.  it  changes  rapidly  into  vapor  without 
rise  of  temperature.  This  vapor,  if  allowed  to  escape  into 
the  atmosphere,  cools  and  condenses  quickly  into  a  cloud 
of  minute  drops  of  water.  This  cloud  is  popularly  called 
steam.  Scientifically,  steam  is  invisible.  What  we  call 
steam  is  a  mass  of  very  small  particles  of  water.  This 
may  be  illustrated  by  boiling  water  in  a  large  flask.  The 
inside  of  the  flask  is  perfectly  transparent,  although  there 
is  a  cloud  of  "  steam  "  issuing  from  its  mouth. 

Water  boils  when  its  vapor  escapes  with  sufficient  pressure  to  over- 
come the  pressure  of  the  atmosphere  upon  its  surface.  Hence  the  boil- 
ing point  depends  upon  the  pressure  —  either  of  the  atmosphere  or  of 
the  vapor  within  the  vessel.  The  boiling  point  is  ioo°C.  (or  2I2°F.) 
when  the  atmospheric  pressure  is  normal,  i.e.  760  mm.  The  boiling 
point  is  lower  as  the  pressure  is  decreased  and  higher  as  the  pressure  is 
increased.  Warm  water  will  boil  under  the  receiver  of  an  air  pump  or 
on  the  top  of  a  high  mountain.  In  the  city  of  Mexico  (7500  feet  above 
sea  level)  water  boils  at  about  92°  C.,  and  in  Quito  in  South  America 
(9350  feet  above  sea  level)  water,  which  boils  at  about  90°  C.,  is  not 
hot  enough  to  cook  potatoes. 

The  pressure  which  water  vapor  exerts  as  it  escapes  from  a  liquid  is 
called  its  vapor  tension.  Since  the  rate  of  evaporation  depends  upon 
the  temperature  of  the  liquid,  vapor  tension  varies  with  the  temperature. 
Vapor  tension  is  usually  expressed  in  millimeters  of  mercury.  Thus,  at 
ioo°C.  the  vapor  tension  of  water  is  760  mm.,  because  at  the  boiling 
point  the  vapor  pressure  is  just  enough  to  overcome  the  opposing 
atmospheric  pressure.  At  20°  C.  the  vapor  tension  of  water  is  17.39  mm- 


Properties  of  Water.  37 

A  liter  of  steajn,  if  it  could  exist  at  o°C.  and  760  mm.  pressure,  would 
weigh  0.806  gm.,  or  nine  times  more  than  a  liter  of  hydrogen. 

Natural  Waters.  —  Water  is  never  found  pure  in  nature. 
Even  rain  water,  which  is  usually  regarded  as  the  purest 
natural  water,  contains  gases  and  dust  washed  from  the 
air.  When  rain  strikes  the  ground  it  begins  at  once  to 
take  up  impurities  from  the  rocks,  soil,  and  vegetation. 
Some  of  the  water  flows  along  the  surface,  becoming  more 
and  more  impure,  and  finally  reaches  the  ocean.  From  25 
to  40  per  cent  of  the  annual  rainfall  in  temperate  regions 
soaks  into  the  ground  and  trickles  through  the  soil  at  an 
estimated  rate  of  0.2  to  20  feet  a  day.  This  underground 
water  finally  finds  its  way  again  to  the  surface  as  a  spring 
or  well,  through  a  lake  or  river,  or  from  a  hillside.  On  its 
journey  underground  the  water  loses  most,  often  all,  of  its 
organic  matter,  —  remnants  of  vegetable  and  animal  matter, 
—  but  dissolves  mineral  matter  and  gases.  If  the  amount 
of  dissolved  matter  in  spring  water  is  large  or  the  kind  of 
matter  is  so  unusual  as  to  give  the  water  a  marked  taste  or 
medicinal  properties,  the  water  is  called  mineral  water. 
Water  containing  calcium  and  magnesium  compounds  is 
hard,  but  in  soft  water,  such  as  rain  water,  these  com- 
pounds are  absent. 

There  are  several  hundred  mineral  springs  in  the  United  States. 
Those  having  a  high  temperature  are  called  thermal,  as  at  Hot  Springs, 
Arkansas,  and  at  Bath,  England.  Many  contain  a  large  proportion  of 
common  salt,  as  at  Saratoga,  New  York.  Others  contain  alkaline  matter 
and  carbon  dioxide  gas,  eg.  Vichy  and  Apollinaris  water.  Sulphur 
springs  contain  solid  or  gaseous  compounds  of  sulphur  —  or  both  —  and 
have  valuable  medicinal  properties.  Some,  like  Hunyadi,  are  bitter; 
but  others,  especially  those  in  New  York  State,  which  contain  gaseous 
sulphur  compounds,  have  a  sweet  taste  but  an  unpleasant  odor.  Cha- 
lybeate waters  contain  soluble  iron  compounds.  Many  waters  contain 
lime  and  magnesium  compounds,  and  a  few  contain  alum.  Most  natural 


Descriptive  Chemistry. 


mineral  waters  contain  traces  of  a  large  number  of  different  substances. 
Many  commercial  mineral  waters  have  doubtful  medicinal  value. 

River  water  obviously  contains  the  impurities  brought 
by  springs  and  the  surface  water ;  it  is  also  often  made 
very  impure  by  decaying  animal  and  vegetable  matter, 
which  has  been  purposely  or  accidentally  introduced,  espe- 
cially if  the  river  passes  through  a  thickly  settled  region. 
A  sluggish  river  is  more  apt  to  be  impure  than  a  swift 
one,  because  the  latter  tends  to  purify  itself  by  exposing 
its  impurities  to  the  oxidizing  power  of  the  air.  Ocean 
water  contains  a  large  proportion  of  common  salt.  The 
other  substances  in  order  of  abundance  are  magnesium 
chloride,  magnesium  sulphate,  calcium  sulphate,  and  potas- 
sium chloride  ;  many  other  substances  are  present  in  small 
quantities.  The  peculiar  taste  of  ocean  water  is  due  to 
the  presence  of  these  substances,  and  since  by  evaporation 
the  water  only  is  removed,  the  ocean  always  has  a  "  salty  " 
taste.  The  composition  of  some  natural  waters  is  sum- 
marized in  the  following  — 

TABLE  OF  COMPOSITION  OF  NATURAL  WATERS. 


SOLIDS  —  PARTS  PER  100,000. 

GASES  —  CUBIC  CENTIME- 
TERS PER  LITER. 

KINDS  OF  WATER. 

Organic 
Matter. 

Calcium 
Com- 

Magne- 
sium 
Com- 

Com- 
mon 

Total 
Residue. 

Nitro- 
gen. 

Oxygen. 

Carbon 
Dioxide. 

pounds. 

pounds. 

Salt. 

Rain  . 

I 



— 

•5 

3-4 

I3-I 

6.4 

i-3 

River(Thames) 

34 

20 

1.8 

2.6 

29 

15 

7-4 

30-3 

Spring     .     . 

Trace 

— 

— 

2 

20 

I5.8 

8.6 

i 

Mineral  (Bath) 

Trace 

137 

23 

34 

236 

4 

2 

29 

Ocean      .     . 

Trace 

140 

530 

2650 

3500 

12.  1 

6 

17 

Properties  of  Water.  39 

Drinking  Water.  —  "  Water  used  for  drinking  should  be  free  from 
visible  suspended  particles,  without  disagreeable  taste  or  smell,  and  not 
capable  of  acquiring  such  by  standing  for  a  day  or  two  in  a  clean,  well- 
closed  vessel.  It  should  also  contain  enough  of  the  gases  derived  from 
the  atmosphere  to  give  it  a  slight  taste  distinguishable  from  the  flatness 
of  boiled  or  distilled  water.  It  should  not  contain  solid  matter  in  solu- 
tion to  the  extent  of  more  than  300  parts  in  a  million,  or  about  3  gm. 
to  10  1.  The  mineral  portion  of  this  solid  matter  should  not  con- 
tain any  poisonous  substance.  As  little  as  possible  of  the  solid  contents 
should  consist  of  organic  matter  —  usually  not  exceeding  15  to  20 
parts  per  million,  or  about  2  gm.  to  100  1.  And  it  is  particularly  de- 
sirable that  decomposing  animal  matter  or  substances  which  give  evi- 
dence of  its  previous  presence  should  be  found,  if  at  all,  as  a  mere  trace. 
Above  all,  drinking  water  should  be  free  from  disease-producing  bacte- 
ria or  other  injurious  microorganisms." 

The  problem  of  obtaining  drinking  water  in  large  quantities  is  usually 
local.  In  some  cities  the  water  is  purified  by  filtering  it  through  a 
layer  of  sand  and  gravel,  an  acre  or  more  in  area  and  several  feet  deep. 
Such  a  filter  removes  bacteria  almost  completely,  though  it  must  be 
frequently  cleaned.  Sometimes  the  water  is  stored  in  a  large  settling 
basin  or  reservoir  and  purified  by  adding  alum,  or  a  similar  substance, 
which  causes  the  suspended  matter  to  settle.  Dissolved  substances 
cannot  be  removed  without  considerable  difficulty,  so  as  a  rule  water 
is  taken  from  a  source  which  is  reasonably  pure. 

The  purity  of  drinking  water  is  usually  determined  by  a  water  analy- 
sis. This  is  not  a  decomposition  of  water,  but  a  chemical  examination 
of  a  sample  for  the  presence  and  amount  of  certain  substances  which 
indicate  or  cause  impurity.  A  chemical  examination  is  of  limited  value, 
however,  unless  it  is  supplemented  by  a  microscopic  study  of  a  fresh 
sample  and  a  rigid  sanitary  inspection  of  the  premises.  Water  which 
is  clear,  sparkling,  cool,  attractive  to  the  eye,  and  pleasant  to  the  taste 
may  be  seriously  polluted  by  disease  germs,  or  may  be  liable  to  sudden 
contamination  from  some  unsuspected  source.  On  the  other  hand,  a 
rather  unpleasant-looking  water  may  be  harmless.  Hence  the  necessity 
of  careful  and  extended  examination  of  water  to  be  used  as  a  beverage. 

Water  may  be  purified  by  distillation.  This  operation  is  not  con- 
venient with  large  quantities.  It  is  performed  in  the  laboratory  in  a 
condenser,  which  is  shown  in  Figure  6  arranged  for  use. 

The  condenser  consists  of  an  outer  tube,  A  A,  provided  with  an  inlet 


4o 


Descriptive  Chemistry. 


and  an  outlet  for  a  current  of  cold  water,  which  surrounds  the  inner 
tube,  BB.    The  vapor  from  the  water  boiling  in  the  flask,  C,  condenses 


FlG.  6.  —  Condenser  arranged  for  the  distillation  of  water. 

in  the  inner  tube,  owing  to  the  decrease  in  temperature,  and  drops  off 
the  lower  end  of  this  tube,  as  the  distillate,  into  the  receiver,  D,  while 
the  impurities  remain  behind  in  the  flask.  Distilled  water  is  prepared 
on  a  large  scale  in  metal  vessels,  and  the  vapor  is  con- 
densed in  a  block  tin  pipe  coiled  around  the  inside  of 
a  vessel  through  which  a 
current  of  cold  water  is  flow- 
ing. This  coiled  pipe  is 
called  a  worm  (Fig.  7). 
Distilled  water  is  used  in 
the  chemical  laboratory ; 
large  quantities  are  made 
into  ice.  Distillation  is  an 
old  process.  A  quaint  still 
is  shown  in  Figure  8.  Dis- 
tillation is  the  process  used 
to  separate  liquids  from 
solids  and  from  each  other, 
FIG.  7.  — Worm-  and  finds  extensive  appli-  FIG.  8.— A  quaint  still, 
shaped  tube.  cation  in  the  manufacture  of  liquors  and  kerosene  oil. 


Properties  of  Water.  41 

Solution.  —  Many  solids,  liquids,  and  gases  disappear 
when  put  into  water.  This  operation  is  called  dissolving, 
or  putting  into  solution.  The  resulting  liquid  is  called  a 
solution  of  the  substance  used.  The  liquid  in  which  the 
substance  dissolves  is  called  the  solvent,  and  the  dissolved 
substance  is  called  the  solute.  If  the  solute  is  not  vola- 
tile, or  not  very  volatile,  it  may  be  recovered  by  evaporat- 
ing, or  distilling  off,  the  water.  The  degree  of  solubility 
is  usually  expressed  by  the  terms  sligJitly  soluble,  soluble, 
and  very  soluble.  It  is  more  accurate,  and  usually  desir- 
able, to  state  the  proportions  of  solvent  and  solute,  and 
also  the  temperature.  Thus,  instead  of  saying  that 
common  salt  is  very  soluble  in  cold  water,  it  is  better  to 
state  that  36  gm.  of  salt  dissolve  in  100  cc.  of  water  at 
20°  C.  Substances  which  do  not  dissolve  in  water  are 
called  insoluble,  though  this  term  is  also  applied  to  those 
substances  a  minute  quantity  of  which  dissolves  in  water. 
Thus  glass,  sand,  and  many  rocks  are  usually  classed  as 
insoluble  substances,  but  they  dissolve  appreciably  in 
water. 

A  solution  which  contains  a  small  proportion  of  solute 
is  called  a  dilute  solution  ;  one  containing  a  large  propor- 
tion is  called  a  concentrated  solution.  Thus,  dilute  sul- 
phuric acid  usually  contains  one  volume  of  acid  to  three 
or  more  volumes  of  water,  while  concentrated  sulphuric 
acid  is  nearly  98  per  cent  acid.  Sometimes  the  terms 
weak  and  strong  replace  dilute  and  concentrated,  but  they 
are  ambiguous,  and  their  use  should  be  avoided. 

Solutions  of  Gases.  — Water  dissolves  or  absorbs  many 
gases.  The  degree  of  solubility  depends  upon  the  gas, 
the  temperature  of  the  water,  and  the  pressure  at  which 
solution  occurs.  Some  gases,  such  as  ammonia  and  hydro- 
chloric acid  gas,  are  very  soluble  in  water.  Advantage  of 


Descriptive  Chemistry. 


this  fact  is  taken  in  manufacturing  ammonium  hydroxide 
and  hydrochloric  acid.  Each  commercial  substance  is 
merely  a  water  solution  of  the  respective  gases,  ammonia 
and  hydrochloric  acid  gas  ;  the  gas  is  readily  liberated  by 
heating  the  liquid. 

The  common  gases,  oxygen  and  hydrogen,  are  only  slightly  soluble 
in  water.  Air  dissolves  in  water,  as  may  easily  be  shown  by  heating 
faucet  water,  bubbles  of  air  forming  and  escaping  quickly  as  heat  is  ap- 
plied. Carbon  dioxide  gas  is  very  soluble  in  water.  Water  containing 
this  gas  is  called  "soda  water,"  or  carbonated  water.  More  gas  is  forced 
into  the  water  than  will  dissolve  at  the  ordinary  temperature  and  pres- 
sure, as  may  be  seen  by  the  rapid  escape  of  gas  when  the  water  is  drawn 
from  a  soda  fountain.  This  rapid  escape  of  a  gas  is  called  efferves- 
cence. "  Soda  water  "  must,  therefore,  be  stored  in  a  strong  vessel  and 
kept  in  a  cool  place.  The  gas  was  formerly  obtained  from  sodium  bi- 
carbonate—  a  compound  related  to  "soda'1;  hence  the  name  "soda 
water.11  It  is  now  prepared  from  marble  and  an  acid,  or  from  liquid 
carbon  dioxide. 

The  volume  of  gas  which  will  dissolve  in  water  decreases  with  rise 
of  temperature.  Thus,  100  cc.  of  water  at  o°C.  will  dissolve  179.6  cc. 
of  carbon  dioxide,  but  only  90.1  cc.  at  20°  C.  The  volume  of  a  mod- 
erately soluble  gas  which  is  dissolved  by  water  is  directly  proportional 
to  the  pressure  if  the  temperature  is  constant.  This  is  Henry's  law. 
It  is  illustrated  by  the  following  — 

TABLE  OF  SOLUBILITY  OF  CARBON  DIOXIDE  GAS. 


VOL.  OF  WATER  AT  o°  C. 

VOL.  OF  CARBON  DIOXIDE    MEASURED 
UNDER  NORMAL  CONDITIONS. 

PRESSURE'  IN 
ATMOSPHERES. 

I  1. 

900  cc. 
1800  cc. 
3600  cc. 

•5 
i 

2 

7200  cc. 

4 

The  tremendous  pressure  to  which  subterranean  gases  are  subjected 
accounts  for  their  presence,  especially  carbon  dioxide,  in  such  large  pro- 
portions in  the  waters  of  mineral  springs. 


Properties  of  Water.  43 

Solutions  of  Liquids.  —  The  solubility  %of  liquids  in 
water  varies  between  wide  limits.  Some,  such  as  alcohol 
and  glycerine,  are  soluble  in  all  proportions.  Oils,  such 
as  kerosene,  are  practically  insoluble;  hence  the  old  adage, 
"  Oil  and  water  will  not  mix."  Carbon  disulphide  is  also 
insoluble,  as  may  be  seen  by  the  formation,  after  agitation, 
of  two  distinct  layers  of  liquid.  The  existence  of  two 
layers,  however,  is  not  always  absolute  proof  of  insolubility. 
JEther  and  water  form  two  layers,  but  each  dissolves  appre- 
ciably in  the  other.  In  many  cases  a  rise  of  temperature 
increases  the  solubility  of  liquids  in  water. 

Solutions  of  Solids.  —  The  solubility  of  solids  in  water 
is  a  matter  of  tremendous  practical  importance.  The 
abundance  of  water  and  its  power  to  dissolve  such  a  vast 
number  of  different  solids  have  led  some  to  call  water 
"  the  universal  solvent."  The  far-reaching  effect  of  this 
marvelous  power  in  nature  and  its  indispensable  value  to 
man  have  been  considered.  (See  above.) 

The  degree  of  solubility  of  solids  in  water  varies  with 
the  substance  and  with  the  temperature  of  the  water. 
Some,  like  potassium  permanganate,  are  very  soluble,  while 
others,  like  calcium  sulphate,  are  difficultly  soluble.  In 
most  cases  solubility  increases  with  a  rise  of  temperature ; 
hence  the  common  practice  of  heating  to  hasten  solution. 
The  effect  of  increased  temperature  on  solubility  is  some- 
times very  marked,  the  solubility  being  increased  fourfold 
in  some  cases.  Calcium  hydroxide  is  less  soluble  in  hot 
than  in  cold  water,  while  common  salt  (sodium  chloride) 
dissolves  to  about  the  same  degree  in  each.  There  is  a 
limit  to  solubility.  That  is,  a  given  volume  of  water  at 
a  fixed  temperature  will  dissolve  a  definite  weight  of  solid 
and  no  more,  although  some  undissolved  solid  remains  in 
the  water. 


44  Descriptive  Chemistry. 

TABLE  OF  SOLUBILITY  OF  SOLIDS  IN  WATER. 


NUMBER  OF  GRAMS  SOLUBLE  IN  100  GRAMS 

SOLIDS. 

OF  WATER  AT 

20°  C. 

IOO°C. 

Calcium  chloride  

74 

155 

Copper  sulphate  (cryst.)     .     . 

42.3 

203.3 

Magnesium  sulphate      .     . 

36.2 

73.8 

Potassium  chlorate   .... 

7.2 

59-5 

Potassium  chloride  .... 

35 

57 

Potassium  dichromate  .     .     . 

13 

102 

Potassium  nitrate      .     .     . 

3'-7 

246 

Potassium  sulphate  .... 

10.6 

26 

Sodium  chloride  

36 

39-7 

A  solution  is  saturated  at  a  given  temperature  when  it 
will  dissolve  no  more  solid.  If  a  hot  solution,  especially 
one  which  contains  much  solid,  is  cooled  slowly,  the  solid 
soon  begins  to  separate  from  the  liquid,  since  solubility 
usually  decreases  with  a  fall  of  temperature.  Often  the 
solid  is  deposited  in  masses  having  a  definite  shape.  This 
operation  is  called  crystallization,  and  the  masses  are 
called  crystals  (see  below).  The  shape  and  color  of  the 
crystal  are  characteristic  of  the  substance,  and  serve  to 
identify  it.  Thus,  common  salt  crystallizes  in  cubes. 
Sometimes  it  is  more  convenient  to  evaporate  a  hot,  con- 
centrated solution.  The  point  of  saturation  at  the  lower 
temperature  is  thus  reached  so  gradually  that  the  crystals 
can  grow  symmetrically.  A  brief  account  of  crystals  will 
be  found  in  §  3  of  the  Appendix. 

A  solid  can  also  be  separated  from  a  solution  by  precipi- 
tation. This  may  be  done  in  two  ways,  (i)  By  adding 
a  liquid  in  which  the  solid  is  not  very  soluble.  Thus,  when 


Properties  of  Water.  45 

water  is  added  to  an  alcoholic  solution  of  camphor,  the 
liquid  becomes  turbid,  or  cloudy,  because  the  camphor  is 
not  soluble  in  water.  That  is,  the  solid  has  been  precipi- 
tated as  very  fine  particles  which  remain  suspended  in  the 
liquid  for  some  time.  Since  the  separated  solid  sooner  or 
later  falls  to  the  bottom  of  the  vessel,  it  is  called  a  precipi- 
tate. (2)  By  changing  the  dissolved  solid  into  another 
substance  not  soluble  in  the  liquid.  Such  chemical  changes 
are  examples  of  double  decomposition.  Thus,  when  so- 
dium chloride  solution  is  added  to  silver  nitrate  solution  a 
white,  curdy  precipitate  of  silver  chloride  is  formed.  A 
soluble  silver  compound  has  thus  been  changed  into  an 
insoluble  silver  compound,  thereby  removing  the  combined 
silver  from  the  solution.  So,  also,  a  soluble  chlorine  com- 
pound (sodium  chloride)  has  been  changed  into  an  insoluble 
chlorine  compound  (silver  chloride),  thereby  removing  the 
combined  chlorine  from  the  solution.  Precipitation  is  a 
very  common  operation  in  chemistry. 

A  hot,  saturated  solution  of  some  solids,  such  as  sodium 
sulphate  and  sodium  thiosulphate,  deposits  no  crystals 
when  the  clear  solution  cools.  Such  solutions  are  super- 
saturated. Supersaturation  can  occur  only  when  the  un- 
dissolved  solid  is  not  present.  Hence,  if  a  fragment  of 
the  solid  is  dropped  into  the  supersaturated  solution,  crys- 
tals soon  begin  to  form  upon  the  fragment,  and  this  sepa- 
ration continues  until  nearly  all  the  substance  is  deposited, 
often  forming  a  solid  mass  in  the  test  tube.  Dust,  or  even 
shaking,  causes  the  substance  to  be  deposited,  hence  the 
solution  should  be  kept  corked  and  left  undisturbed.  Sat- 
uration is  analogous  to  stable  equilibrium,  while  supersatu- 
ration  resembles  unstable  equilibrium. 

Water  of  Crystallization.  —  Crystals  deposited  from 
the  water  solution  of  many  solids,  even  after  they  are  dried 


4.6  Descriptive  Chemistry. 

by  pressing  between  filter  paper  or  by  exposure  to  a  mod- 
erate temperature,  often  contain  water  which  seems  to  be 
an  essential  part  of  the  chemical  compound.  This  water 
is  called  water  of  crystallization.  The  crystals  of  some 
compounds,  e.g.  sodium  carbonate  and  sodium  sulphate, 
lose  their  water  of  crystallization  and  crumble  on  exposure 
to  the  air.  This  property  is  called  efflorescence,  and  such 
crystals  are  said  to  effloresce  or  to  be  efflorescent.  Heat 
will  drive  the  water  of  crystallization  from  crystals  which 
contain  it,  e.g.  gypsum,  alum,  and  copper  sulphate. 

The  proportion  of  water  of  crystallization  in  crystals  is  not  arbitrary. 
It  is  constant  in  the  same  compound  when  crystallized  under  uniform 
conditions,  but  the  proportion  varies  between  wide  limits  in  different 
substances.  No  explanation  has  been  given  of  the  varying  amount  of 
water  of  crystallization,  nor  of  its  necessity  for  the  form  and  color  of 
some  crystals  and  not  for  others.  Some  well-crystallized  substances 
contain  no  water  of  crystallization,  e.g.  potassium  nitrate,  potassium 
dichromate,  sugar,  and  salt. 

Crystals  which  have  lost  their  water  of  crystallization  are  said  to  be 
dehydrated  or  anhydrous.  Thus,  the  grayish  powder  obtained  by 
heating  the  blue  crystallized  copper  sulphate  is  called  dehydrated  cop- 
per sulphate.  The  words  dehydrated  and  anhydrous  have  been  extended 
to  describe  any  substance  from  which  water  has  been  removed,  as  anhy- 
drous alcohol  or  ether.  The  opposite  term,  hydrated,  is  sometimes 
applied  to  a  compound  to  emphasize  the  fact  that  it  contains  water  of 
crystallization. 

Deliquescence.  —  Many  substances,  crystallized  and 
uncrystallized,  absorb  water  when  exposed  to  the  air,  and 
become  moist,  or  even  dissolve  in  the  water.  Calcium 
chloride,  potassium  carbonate,  zinc  chloride,  sodium  hydrox- 
ide, and  potassium  hydroxide  belong  to  this  class.  This 
property  is  called  deliquescence,  and  the  substances  are 
said  to  deliquesce,  or  to  be  deliquescent.  The  term  hygro- 
scopic is  applied  to  substances  which  absorb  water,  but 
hygroscopic  substances  do  not  dissolve  in  the  absorbed 


Properties  of  Water.  47 

water,  and  sometimes  do  not  even  become  moist.     Quick- 
lime is  hygroscopic. 

Common  salt,  or  sodium  chloride,  often  appears  to  deliquesce,  espe- 
cially in  damp  weather.  The  deliquescence  is  due,  however,  to  the 
presence  of  magnesium  and  calcium  chlorides.  Sodium  nitrate  is  some- 
what deliquescent,  and  cannot  be  used  in  the  manufacture  of  gunpowder, 
so  potassium  nitrate  is  used  instead.  This  property  of  deliquescence  is 
often  utilized  in  the  laboratory  to  remove  water  vapor  from  gases,  cal- 
cium chloride  being  especially  serviceable  for  this  purpose. 

Thermal  Phenomena  of  Solution.  —  Solution  is  often  accompanied 
by  an  appreciable  change  of  temperature.  When  sulphuric  acid  is  poured 
into  water,  heat  is  produced.  With  large  quantities  the  heat  is  so  great 
that  the  mixture  often  boils,  and  sometimes  the  hot  acid  is  spattered. 
Hence,  the  acid  should  be  added  slowly  to  the  water,  and  the  mixture 
constantly  stirred.  Other  substances  which  dissolve  with  the  liberation 
of  heat  are  fused  calcium  chloride,  potassium  hydroxide,  and  sodium 
hydroxide.  Some  which  dissolve  with  a  fall  of  temperature  are  crystal- 
lized calcium  chloride,  ammonium  nitrate,  ammonium  chloride,  and 
potassium  nitrate.  This  subject  is  still  under  investigation. 

Solution  and  Chemical  Action.  —  Probably  when  a  sub- 
stance dissolves  it  is  so  modified  that  it  can  participate 
more  readily  in  chemical  changes.  Hence,  solution  is  an 
aid  to  chemical  change,  and  is  often  an  easy  means  of 
causing  it.  Thus,  if  dry  tartaric  acid  and  sodium  bicar- 
bonate are  mixed,  there  is  no  evidence  of  chemical  action ; 
but  when  the  mixture  is  poured  into  water,  the  copious 
evolution  of  carbon  dioxide  gas  is  conclusive  evidence  of  a 
chemical  change.  Similarly,  when  a  dry  mixture  of  ferrous 
sulphate  and  potassium  ferrocyanide  is  poured  into  water, 
the  immediate  appearance  of  a  blue  precipitate  shows  that 
the  water  was  needed  for  the  chemical  change.  Solution 
is  such  an  important  aid  to  chemical  action  that  many 
substances  employed  in  the  laboratory  are  in  solution, 
and  many  processes  in  chemistry  are  "  wet "  processes. 


48  Descriptive  Chemistry. 

Mention  has  already  been  made  of  the  application  of  this 
fact  to  many  industries. 

The  Nature  of  Solution  has  long  been  a  subject  of  specu- 
lation and  study.  The  problem  as  a  whole  is  still  unsolved, 
though  much  light  has  been  thrown  upon  the  question  by 
recent  investigations  (see  Chapter  X). 

EXERCISES. 

1 .  Mention  several  familiar  properties  of  water. 

2.  In  what  forms  does  water  exist  ? 

3.  Give  the  per  cent  of  water  in  some  familiar  foods. 

4.  Develop  the  topics  :    («)  water  is  an  erosive  agent ;    (£)  water 
is  a  solvent  in  nature ;    (V)  water  has  many  industrial  applications ; 

(d)  water  behaves  exceptionally  when  heated  from  o°  C.  to  io°C.; 

(e)  ice  floats  ;  (/)  water  is  a  cleansing  agent. 

5.  Explain  these  expressions :  (a)  water  has  its  maximum  density 
at  4°  C. ;  (b)  the  density  of  ice  is  0.92  ;  (V)  steam  is  invisible  ;   (d)  the 
lower  the  pressure,  the  lower  the  boiling  point ;  (e)   10  cc.  per  liter ; 
(_/")  parts  per  million. 

6.  How  do  natural  waters  illustrate  the  solvent  power  of  water? 

7.  What  is  (a)   mineral  water,    (£)   soft  water,    (c)   hard  water, 
(d)  sulphur  water,  (e}  chalybeate  water  ? 

8.  What  does  ocean  water  contain?     Why  is  the  sea  water  salt? 

9.  What  constitutes  a  safe  drinking  water?     How  may  city  water 
be  purified  ?     What  is  a  water  analysis  ? 

10.  Describe  the  operation  of  distillation.     What  is  a   condenser 
and  why  is  it" so  named?     Is  distillation  a  new  or  an  old  process? 
Of  what  industrial  use  is  it? 

11.  Define  and  illustrate  (a)  water  of  crystallization,    (b)  efflores- 
cence, (c}  deliquescence,    (d)  hygroscopic,  (e)  anhydrous,  (/)  dehy- 
drated, (g)  crystal,  (//)  crystallization. 

12.  Define   and   illustrate    (a}   solution,    (b}    solvent,    (V)    solute, 
(d)  soluble,  (e)  slightly  soluble,  (/)  very  soluble,  (g}  insoluble,  (h}  di- 
lute,   (/)    concentrated,    (/)    saturated    solution,    (£)    supersaturated 
solution. 

13.  Give  several  facts  about  the  solubility  of  gases  in  water.     What 
is  (a)  soda  water,  (b}  carbonated  water?     How  do  we  know  that  air 


Properties  of  Water.  49 

dissolves  in  water?  Why  do  subterranean  waters  often  contain  dis- 
solved gases?  State  Henry's  law  of  the  solubility  of  gases.  What 
effect  has  (a)  heat  and  (£)  cold  on  the  solubility  of  gases  in  water? 

14.  What  liquids  are  soluble  in  water?     How  may  such  liquids  be 
separated  from  water? 

15.  What  general  effect  has  (a}  heat  and  (£)  cold  on  the  solubility 
of  solids  in  water?     Mention  some  solids  which  are  (a)  very  soluble, 
(b)  moderately  soluble,    (<:)  almost  insoluble  in  water.     Develop  the 
topic  :  There  is  a  limit  to  the  solubility  of  solids  in  water. 

1 6.  (a}   How  would  you  find  the  approximate  amount  of  water  in 
(i)  milk,  (2)  an  apple?     (b)  How  would  you  find  the  per  cent  of  each 
substance  in  a  mixture  of  sand  and  sugar? 

17.  Develop  the  topic:  Solution  aids  chemical  change.     Why  are 
so  many  solutions  used  in  a  laboratory  ? 

1 8.  What  changes  in  volume  occur  when  (a}  ice  melts,  (£)  water 
freezes,  (c)  water  is  heated  from  o°C.  to  i5°C.,  (d)  water  is  cooled 
from  I5°C.  to  o°C.? 

19.  Write  an  essay  on  "Mineral  Springs  in  the  United  States." 


PROBLEMS. 

1.  If  1.5  gm.  of  crystallized  barium  chloride  lose  0.22  gm.  when 
heated  to  constant  weight,  what  per  cent  of  water  of  crystallization 
does  it  contain? 

2.  If  2   gm.  of  another  lot  of  barium  chloride   lose  0.295  gm., 
what  per  cent  of  it  was  water  of  crystallization  ? 

3.  If  a  liter  of  sea  water  has  a  density  of  1.25,  how  many  grams  of 
"salt"  does  it  contain? 

4.  If  the  density  of  ice  is  0.92,  what  volume  will  a  liter  of  water  at 
4°C.  occupy  when  frozen?  Ans.  1.087  !• 

5.  How  much  water  (approximately)  is  contained  in  (a)  2  Ib.  of 
lobster,  (b}  56  Ib.  of  potatoes,  (c)  i  Ib.  of  tomatoes,  {d}  2  Ib.  of  milk, 
(e)  i  Ib.  of  white  bread,  (/)  a  human  body  weighing  150  Ib.  ? 

6.  If  a  kilogram  of  sea  water  contains  36.4  gm.  of  "salt,"  what 
per  cent  of  the  water  is  "  salt "  ? 

7.  If  a  block  of  ice  weighs  280  kg.,  what  is  its  volume? 

Ans.  304.3  1. 

8.  A  solution  measures  100  cc.  and  contains  15  gm.  of  potassium 
nitrate.     What  per  cent  of  water  and  of  solid  is  in  the  solution  ? 


CHAPTER   V. 
COMPOSITION  OF  WATER. 

-  WATER  was  considered  an  element  until  about  the  end 
of  the  eighteenth  century.  At  that  time  it  was  shown  to 
be  a  compound  of  hydrogen  and  oxygen.  Many  famous 
chemists  worked  on  this  problem. 

The  Composition  of  a  Compound  is  determined  either 
by  analysis  or  synthesis,  i.e.  by  taking  it  apart  or  putting 
its  parts  together.  Sometimes  both  methods  are  used, 
since  each  method  fortifies  the  other  and  strengthens  the 
final  conclusion.  These  methods  find  excellent  application 
in  determining  the  composition  of  water. 

Analysis  and  synthesis  may  be  qualitative  or  quantitative.  A  quali- 
tative experiment  is  a  study  of  the  properties  of  elements  and  com- 
pounds with  a  view  of  discovering  what  they  contain.  A  quantitative 
experiment  is  an  accurate  determination  of  the  weight  or  volume  of  the 
components  of  a  compound.  Qualitative  tests  involve  merely  quality, 
while  in  quantitative  tests  quantity  is  the  essential  feature.  Obviously, 
a  complete  determination  of  the  composition  of  a  compound  requires 
both  tests. 

Water  contains  Hydrogen.  —  When  steam  is  passed 
over  heated  metals,  hydrogen  is  liberated.  Lavoisier's 
demonstration  of  this  fact  has  already  been  considered 
(see  Preparation  of  Hydrogen).  The  fact  that  sodium 
liberates  hydrogen  from  water  at  the  ordinary  temperature 
has  also  been  discussed  (see  ibid.).  If  red  litmus  paper  is 
put  into  the  water  from  which  the  sodium  has  liberated 
hydrogen,  the  litmus  paper  becomes  blue.  This  change 

50 


Composition  of  Water.  51 

i 

of  color  from  red  to  blue  shows  that  an  alkali  is  in  the 
water,  because  alkalies  turn  red  litmus  paper  blue.  The 
alkali  is  sodium  hydroxide,  and  it  may  be  obtained  as  a 
white  solid  by  evaporating  the  water.  Sodium  hydroxide 
is  a  compound  of  sodium,  hydrogen,  and  oxygen,  and  is 
formed  by  replacing  part  of  the  hydrogen  of  water  by 
sodium.  Since  sodium  liberates  hydrogen  from  water,  and 
forms  at  the  same  time  a  compound  —  sodium  hydroxide 
—  containing  hydrogen,  the  hydrogen  in  water  must  be 
divisible  into  two  parts.  Now  if  o.  I  gm.  of  sodium  is 
allowed  to  act  upon  water,  48.22  cc.  of  hydrogen  are  liber- 
ated ;  and  if  the  sodiunThydroxide  thus  formed  is  dried  and 
heated  with  sodium,  48.22  cc.  more  of  hydrogen  are  ob- 
tained. This  shows  that  the  hydrogen  in  water  is  divisible 
into  two  equal  parts  —  a  fact  which  will  soon  be  utilized. 

Water  contains  Oxygen.  —  The  fact  that  oxygen  is  a 
component  of  water  has  already  been  suggested,  e.g.  (i) 
by  the  production  of  water  when  hydrogen  is  burned  in 
air,  (2)  by  the  formation  of  a  compound  of  iron  and  oxy- 
gen when  steam  is  passed  over  hot  iron,  and  (3)  by  the 
formation  of  sodium  hydroxide  when  sodium  acts  upon 
water.  These  proofs,  however,  are  all  indirect.  A  simple 
direct  demonstration  of  the  presence  of  oxygen  in  water 
may  be  made  by  allowing  chlorine  water  to  stand  in  the 
sunlight.  (Chlorine  water  is  prepared  by  saturating  water 
with  chlorine  gas  —  an  element  to  be  studied  in  Chapter 

XI.)     A  long  tube  like  that  shown        ^ 

in    Figure    9    is    completely    filled    GU — 

with  chlorine  water,  the  open  end  is    FIG.  9.  —  Tube  for  decompo- 

,  L    .     .  sition  of  water  by  chlorine. 

immersed    in    a    vessel     containing 

some  of  the*  same  solution,  and  the  whole  apparatus  is 
placed  in  the  direct  sunlight.  Bubbles  of  gas  soon  appear 
in  the  liquid,  and  after  a  few  hours  a  small  volume  of 


Descriptive  Chemistry. 


gas  collects  at  the  top  of  the  tube.      This  gas  may  be 
shown,  by  the  usual  tests,  to  be  oxygen. 

The  Electrolysis  of  Water  is  its  decomposition  by  elec- 
tricity.    It   is   accomplished   in   the  apparatus  shown  in 

Figure  i  o.  Since  pure  water  does 
not  conduct  electricity,  sulphuric 
acid  is  added.  Enough  of  this 
acid  mixture  is  poured  into  the 
apparatus  to  fill  the  reservoir 
half  full  after  the  stopcocks  have 
been  closed.  As  soon  as  an 
electric  battery  of  two  or  more 
cells  is  connected  by  wires  with 
the  piece  of  platinum  near  the 
bottom  of  each  tube,  bubbles 
of  gas  form  on  the  platinum, 
and  as  the  action  proceeds,  the 
bubbles  rise  and  displace  the 
water  in  each  tube.  The  volume 
of  gas  is  greater  in  one  tube. 
Assuming  that  the  tubes  have 
the  same  diameter,  the  volumes 
are  in  the  same  ratio  as  their 
heights,  which  will  be  found  by 
measurement  to  be  two  to  one. 
The  larger  volume  of  gas  is 
FIG.  io.  — Hofmann  apparatus  for  hydrogen  and  the  smaller  one 

electrolysis  of  water. 

is  oxygen.  Many  accurate  repe- 
titions of  this  experiment  have  shown  that  only  hydrogen 
and  oxygen  are  produced,  and  that  the  ratio  of  their  volumes 
is  two  to  one.  It  has  also  been  shown  that  the  sum  of  the 
weights  of  the  two  gases  equals  the  weight  of  the  water 
decomposed.  The  whole  experiment  demonstrates  that 


Composition  of  Water.  53 

water  is  a  compound  consisting  of  two  volumes  of  hydro- 
gen combined  with  one  volume  of  oxygen. 

Water  was  first  decomposed  by  electricity  in  1800  by  Nicholson  and 
Carlisle.  Davy  confirmed  their  work  by  a  series  of  brilliant  experi- 
ments extending  through  a  period  of  six  years  (1800-1806).  During 
this  time  he  not  only  proved  that  the  volume  of  hydrogen  is  double  that 
of  oxygen,  but  by  electrolyzing  water  in  a  gold  vessel  placed  in  an  atmos- 
phere of  hydrogen,  he  proved  that  nothing  but  these  gases  is  produced. 

The  Quantitative  Composition  of  Water.  —  The  fore- 
going facts  about  the  composition  of  water  have  been 
mainly  qualitative.  They  have  shown  by  analysis  and 
synthesis  that  water  consists  of  hydrogen  and  oxygen,  and 
that  the  ratio  of  their  volumes  is  approximately  two  to  one. 
Decisive  evidence  of  the  quantitative  composition  of  water 
is  obtained  by  a  determination  of  its  volumetric  and  its 
gravimetric  composition.  Volumetric  means  "by  volume" 
and  gravimetric  means  "  by  weight." 

The  Volumetric  Composition  of  Water  is  determined 
by  exploding  a  mixture  of  known  volumes  of  hydrogen 
and  oxygen  in  a  eudiometer. 

Gas  volumes  which  are  to  be  compared  with  each  other  must  be  dry 
and  at  the  same  temperature  and  pressure.  This  requirement,  which  is 
called  the  "  standard  condition,"  is  inconvenient,  and  almost  impracti- 
cable. Hence,  it  is  customary  to  measure  each  volume  of  moist  gas 
under  the  existing  conditions,  and  then  reduce  the  observed  volume  to 
that  volume  which  the  gas  would  occupy  if  standard  conditions  pre- 
vailed. The  reduction  to  standard  conditions  is  accomplished  by  the 
formula  —  j/r  /pi  _  n\ 


760(1  +  .  00366  /) 

In  the  formula l  —    V  =  the  corrected  volume. 
V  =  the  observed  volume. 


1  A  complete  discussion  of  the  laws  of  gases,  the  principles  which  control  their 
measurement,  together  with  the  development  of  the  above  formula  for  reduction  to 
standard  conditions,  may  be  found  in  Appendix  B  of  the  author's  "  Experimental 
Chemistry."  See  also  the  Laws  of  Boyle  and  Charles  in  Chapter  II,  and  Vapor 
Density  in  Chapter  IV  (this  book). 


54 


Descriptive  Chemistry. 


Pf=  the  observed  pressure. 
/    =  the  observed  temperature. 
a    —  the  vapor  tension  at  /°  C. 

A  convenient  form  of  apparatus  for  determining  the  volu- 
metric composition  of  water  is  shown  in  Figure  n.  The 
essential  part  is  the  eudiometer,  F.  In  this  graduated 
glass  tube  the  gases  are  accurately  measured  and  ex- 
ploded. The  electric  spark  which  causes  the  explosion  is 

obtained  from  an  induc- 
tion coil  and  battery. 
The  spark  leaps  across 
the  space  between  the 
platinum  wires  at  the 
top  of  the  eudiometer, 
and  the  heat  produced 
by  this  spark  causes  the 
hydrogen  and  oxygen 
to  combine  and  form 
water.  Oxygen  and  hy- 
drogen are  introduced 
separately  into  the  eudi- 
ometer, measured,  and 
-  exploded.  After  the 

FIG.  ii.  — Apparatus  for  determining  the  volu-     explosion,  which  IS  indi- 

cated  by  a  slight  click 

or  flash  of  light,  water  from  the  reservoir,  E,  rushes  up 
into  the  eudiometer.  The  water  does  not  completely  fill 
the  tube,  because  an  excess  of  one  gas  is  added.  This 
additional  gas  takes  no  part  in  the  chemical  change,  but 
merely  serves  to  lessen  the  violence  of  the  explosion,  which 
otherwise  might  break  the  eudiometer.  The  quantity  of 
water  formed  by  the  union  of  the  hydrogen  and  oxygen 


Composition  of  Water. 


55 


is  too  minute  to  measure.  Repeated  trials  of  this  experi- 
ment show  that  two  volumes  of  hydrogen  always  combine 
with  one  volume  of  oxygen.  This  is  the  volumetric  com- 
position of  water. 

The  discovery  of  the  volumetric  composition  of  water  was  not 

made  by  one  chemist  alone.  Priestley,  about  1780,  noticed  that  when 
a  mixture  of  air  and  hydrogen  was  exploded,  "  the  inside  of  the  glass, 
though  clear  and  dry  before,  immediately  became  dewy."  Cavendish, 
in  1781,  showed  that  when  a  mixture  of  two  parts  hydrogen  and  one 
part  oxygen  was  exploded,  nothing  but  water  was  formed.  Watt,  in 
1783,  was  the  first  to  state  that  water  is  a  compound,  though  he  per- 
formed no  experiments  and  probably  did  not  understand  the  real  nature 
of  its  components.  Lavoisier  in  the  same  year  verified  many  facts  pre- 
viously noticed  but  not  completely  understood,  and  undoubtedly  first 
clearly  recognized  and  stated  what  his  contemporaries  had  overlooked. 
The  final  proof  of  the  volumetric  composition  of  water  was  an  accurate 
verification  in  1805  by  Gay-Lussac  and  Humboldt  of  the  previous  ob- 
servation that  two  volumes  of  hydrogen  unite  with  one  volume  of  oxygen. 

The  Gravimetric  Composition  of  Water  is  determined 
by  passing  dry  hydrogen  over  copper  oxide.  The  method 
depends  upon  the  fact  that  many  oxides,  such  as  those  of 
lead,  copper,  and  iron,  when  heated  in  a  current  of  hydro- 


=7)        (e=  =  l 

i 

i 

c                     c' 

B 

FIG.  12.  —  Apparatus  for  determining  the  gravimetric  composition  of  water. 

gen,  give  up  their  oxygen,  or,  chemically  speaking,  these 
oxides  are  reduced  to  metals.  By  this  reduction  the  oxy- 
gen of  the  oxide  combines  with  the  hydrogen,  thereby 
forming  water  which  is  collected  in  a  weighed  tube. 


56  Descriptive  Chemistry. 

A  convenient  form  of  apparatus  is  shown  in  Figure  12. 
The  copper  oxide  is  placed  in  the  combustion  tube,  CC, 
which  is  made  of  hard  glass.  The  Marchand  tube,  D, 
which  is  filled  with  calcium  chloride,  collects  and  retains 
the  water  formed  in  the  combustion  tube,  as  the  hydrogen 
passes  over  the  hot  copper  oxide.  The  tubes  A,  B,  and  E 
keep  moisture  out  of  the  apparatus.  The  experiment  is 
very  simple.  Copper  oxide  is  placed  in  the  combustion 
tube,  which  is  then  carefully  weighed.  The  Marchand 
tube,  being  filled  with  calcium  chloride,  is  also  weighed. 
After  the  other  tubes  are  properly  filled  and  the  hydrogen 
generator  adjusted,  the  tubes  are  connected  as  shown  in 
the  figure.  The  combustion  tube  is  now  heated,  and  mois- 
ture collects  in  it;  as  the  heat  increases  the  copper  oxide 
glows,  and  the  moisture  passes  into  the  Marchand  tube. 
When  the  operation  is  over  and  the  apparatus  is  cool  and 
free  from  hydrogen,  the  combustion  tube  and  Marchand 
tube  are  weighed.  The  gain  in  weight  of  the  Marchand 
tube  is  the  weight  of  the  water  formed,  while  the  loss  in 
weight  of  the  combustion  tube  is  the  weight  of  the  oxygen 
contained  in  this  water.  An  illustration  will  make  this 
clear.  Dumas  and  Stas,  who  performed  this  experiment 
accurately  in  1843,  found  substantially  that  the  combus- 
tion tube  lost  5.251  gm.  of  oxygen,  while  the  Marchand 
tube  gained  5.909  gm.  of  water.  But  5.251  and  5.909 
are  in  the  same  ratio  as  8  and  9.  Thus  :  — 

5.251  :  5.909  :  :  8  :  9. 

This  means  that  oxygen  makes  up  f  of  water.  The  re- 
maining ^  is  of  course  hydrogen.  In  other  words,  the 
gravimetric  composition  of  water  is  eight  parts  oxygen 
and  one  part  hydrogen.  This  ratio  is  often  stated  in  per- 
centage;  thus  water  contains  — 


Composition  of  Water.  57 

88.88  per  cent  of  oxygen. 
1 1 . 1 1  per  cent  of  hydrogen. 

For  reasons  which  will  soon  be  given,  it  is  more  conven- 
ient to  state  the  composition  of  water  by  weight,  as  two 
parts  hydrogen  to  sixteen  parts  oxygen. 

The  gravimetric  composition  of  water  was  first  determined  about 
1820  by  Berzelius  and  Dulong.  Their  work  was  verified  by  Dumas 
and  Stas  in  1843. 

A  Comparison  of  the  Volumetric  and  Gravimetric  Com- 
position of  Water  shows  that  the  results  of  the  two 
methods  agree.  One  volume  of  oxygen  is  sixteen  times 
heavier  than  an  equal  volume  of  hydrogen  (see  Density  of 
Hydrogen).  Therefore,  the  one  volume  of  oxygen  must 
be  eight  times  heavier  than  the  two  volumes  of  hydrogen 
in  water.  That  is,  the  oxygen  in  water  weighs  eight  times 
more  than  the  hydrogen.  But  this  is  the  ratio  actually 
found  in  determining  the  gravimetric  composition  of  water 
by  an  independent  experiment.  These  facts  strengthen 
our  belief  that  the  composition  of  water  is  — 

By  weight,  one  part  hydrogen  and  eight  parts  oxygen. 
By  volume,  two  parts  hydrogen  and  one  part  oxygen. 

Summary.  — The  following  facts  have  been  shown  con- 
cerning the  composition  of  water :  — 

(1)  Water  is  a  chemical  compound  'of  hydrogen  and 
oxygen. 

(2)  It  is  formed  when  hydrogen  is  burned  in  air,  or 
when  a  mixture  of  hydrogen  and  oxygen  is  exploded. 

(3)  It  can  be  decomposed  by  electricity  into  hydrogen 
and  oxygen  in  the  ratio  of  two  volumes  of  hydrogen  to  one 
volume  of  oxygen. 


58  Descriptive  Chemistry. 

(4)  Sodium  liberates  hydrogen  from  water  and  forms  at 
the  same  time  a  solid  containing  a  quantity  of  hydrogen 
equal  to  the  quantity  of  hydrogen  liberated.     Iron,  other 
metals,  and  carbon  liberate  hydrogen  from  water,  forming 
at  the  same  time  an  oxide  of  the  respective  substance. 

(5)  Chlorine  liberates  oxygen  from  water. 

(6)  Two  volumes  of  hydrogen,  when  exploded  with  one 
volume  of  oxygen,  combine  to  form  water,  and  the  weight 
of  the  water  formed  equals  the  weight  of  the  gases  used. 

(7)  Water  is  formed  by  the  union  of  two  parts  by  weight 
of  hydrogen  and  sixteen  parts  by  weight  of  oxygen. 

EXERCISES. 

1.  How  is  the  composition  of  a  compound  determined  ? 

2.  Define  (a)  synthesis,  (£)  analysis,  (<:)  qualitative,  (//)  quantita- 
tive, (e)  volumetric,  (/)  gravimetric. 

3.  How  would  you  prove  that  water  is  composed  of  hydrogen  and 
oxygen  ? 

4.  How  do  we  know  that  the  hydrogen  in  water  is  divisible  into  two 
equal  parts  ? 

5.  What   is  the  electrolysis  of  water  ?     How  is  it  accomplished  ? 
What  does  it  prove  about  the  composition  of  water  ?     When  and  by 
whom  was  it  first  performed  ?     What  did  Davy  contribute  toward  the 
solution  of  the  problem  ? 

6.  What  is  the  volumetric  composition  of  water  ?     How  is  it  deter- 
mined ?     Who  worked  on  this  problem,  and  what  did  each  contribute 
to  its  solution  ? 

7.  Answer  the  same  questions  (as  in  6)  about  the  gravimetric  com- 
position of  water. 

8.  Compare  the  volumetric  and   the  gravimetric   composition  of 
water. 

9.  What  does  the  burning  of  hydrogen  show  about  the  composition 
of  water  ? 

10.    Summarize   the   essential   facts   regarding  the  composition  of 
water. 


Composition  of  Water.  59 

ii.  Give  a  brief  biographical  account  of  (a)  Nicholson  and  Carlisle, 
(£)  Dumas,  (c}  Humboldt,  (d}  Stas,  (e)  Watt,  (/)  Gay-Lussac  (see 
Appendix,  §  4) . 

PROBLEMS. 

1.  What  weight  of  (a)  hydrogen  and  (fr)  oxygen  can  be  obtained 
by  decomposing  125  gm.  of  water  ? 

2.  What  volume  of  (a)  hydrogen  and  (6)  oxygen  can  be  obtained 
by  decomposing  9  1.  of  water  ? 

3.  What  weight  of  hydrogen  must  unite  with  16  gm.  of  oxygen  to 
form  water  ?     What  weight  with  (#)  40  gm.,  (b)  70  gm.,  (c)  160  gm.  ? 

4.  What  volume  of  oxygen  must  unite  with  2  1.  of  hydrogen  to  form 
water  ?    What  volume  with  (a)  40  1.,  (£)  40  cc.,  (c)  40  qt,  (d )  95  vol- 
umes, (e)  1 60  1.  ? 

5.  What  volume  of  oxygen  is  necessary  to  unite  with  100  gm.  of 
hydrogen  to  form  water  ?     (Suggestion  :  What  is  the  weight  of  a  liter 
of  oxygen  ?) 

6.  Hydrogen  is  passed  over  2.48  gm.  of  hot  copper  oxide,  which  at 
the  end  of  the  experiment  weighed  2.24  gm. ;  the  water  formed  weighed 
0.27  gm.     In  what  ratio  did  the  hydrogen  and  oxygen  combine  ? 

7.  Berzelius  and  Dulong,  in  1820,  obtained  the  following  results  in 
their  determinations  of  the  gravimetric  composition  of  water :   Loss  of 
weight  of  copper  oxide  (in  grams),  10.832  and  8.246.    Weight  of  water 
formed,  12.197  and  9.27.      Calculate  in  each  case  the  ratio  in  which  the 
hydrogen  and  oxygen  combined.     What  is  the  average  ratio  ? 

8.  Dumas  and  Stas  repeated  the  above  work  in  1843,  and  found  as 
an  average  of  nineteen  determinations,  that  840.161  gm.   of  oxygen 
formed  945.439  gm.  of  water.     Calculate  the  ratio  of  combination. 

Hydrogen  Dioxide  is  a  liquid  composed  of  hydrogen  and  oxygen. 
But  the  proportion  of  the  components  is  not  the  same  as  in  water.  It 
contains  two  parts  of  hydrogen  and  thirty-two  parts  of  oxygen  by 
weight.  It  is  often  called,  especially  in  commerce,  hydrogen  peroxide, 
because  its  relative  proportion  of  oxygen  is  greater  than  in  water  —  the 
other  hydrogen  oxide. 

It  is  manufactured  by  treating  barium  dioxide  (or  peroxide)  with 
sulphuric  or  hydrochloric  acid.  The  commercial  solution  has  a  vari- 
able strength,  and  usually  contains  three  or  more  per  cent  of  hydrogen 
dioxide.  It  has  a  sharp,  pungent  odor,  and  a  bitter,  metallic  taste. 


60  Descriptive  Chemistry. 

Hydrogen  dioxide  is  an  unstable  compound ;  it  decomposes  slowly  at 
the  ordinary  temperature,  and  very  rapidly  if  heated.  The  dilute,  com- 
mercial solution  is  somewhat  stable,  but  heat  decomposes  it  completely 
into  water  and  oxygen.  The  ease  with  which  it  yields  oxygen  makes 
it  a  good  oxidizing  agent.  In  this  respect,  hydrogen  dioxide  resembles 
ozone,  and,  indeed,  they  are  sometimes  mistaken  for  each  other.  It  is 
also  a  reducing  agent,  and  is  frequently  used  as  such  in  the  laboratory. 
It  is  used  extensively  to  bleach  animal  and  vegetable  matter,  such  as 
human  hair,  ostrich  feathers,  fur,  silk,  wool,  cotton,  bone,  and  ivory.  It 
is  also  used  as  an  antiseptic  and  disinfectant  in  surgery.  Large  quanti- 
ties are  used  to  restore  the  color  to  faded  paintings  —  a  use  suggested 
by  The'nard,  the  discoverer.  In  the  laboratory  it  is  proving  a  service- 
able reagent. 

Hydrogen  dioxide  is  found  in  the  air,  in  rain  and  snow,  but  the 
proportion  is  variable  and  exceedingly  small. 


CHAPTER   VI. 
THE  ATMOSPHERE  — NITROGEN. 

The  Atmosphere  is  the  great  mass  of  gas  surrounding 
the  earth  and  extending  into  space.  Its  estimated  height 
is  fifty  to  several  hundred  miles.  We  live  at  the  bottom 
of  this  vast  ocean  of  air,  as  it  is  often  called. 

Aristotle  (384-322  B.C.)  regarded  air  as  one  of  the  four  elementary 
principles  whose  combinations  made  up  all  substances  in  the  universe. 
The  other  three  were  earth,  fire,  and  water.  He  taught  that  air  pos- 
sesses two  fundamental  properties, — heat  and  dampness.  The  early 
chemists  used  the  word  air  in  the  sense  in  which  the  word  gas  is  now 
employed.  Thus,  we  have  already  learned  that  hydrogen  was  first 
called  inflammable  air. 

The  terms  atmosphere  and  air  are  often  used  inter- 
changeably, though  by  air  we  usually  mean  a  limited  por- 
tion of  the  atmosphere.  Many  skillful  chemists  have 
studied  the  action  of  air  on  living  things,  its  relation  to 
combustion,  the  effect  of  its  weight,  its  composition,  and 
its  varied  properties.  Their  work  has  contributed  many 
fundamental  facts  to  science. 

General   Properties  of   the   Atmosphere.  —  Air    has 

weight.  We  often  use  the  expression  "  light  as  air."  But 
a  cubic  foot  of  air  weighs  1.28  oz.  and  a  room  40  x  50  X  25 
ft.  contains  about  two  tons  of  air.  The  total  weight  of  the 
atmosphere  has  been  estimated  to  be  five  thousand  millions 
of  millions  of  tons.  This  enormous  mass  resting  upon  the 
earth  exerts  a  pressure  which  is  about  fifteen  pounds  on 
every  square  inch.  This  amount  of  pressure  upon  a 

61 


62  Descriptive  Chemistry. 

square  inch  is  called  "an  atmosphere,"  and  it  is  some- 
times used  as  a  unit  of  pressure.  Thus,  three  atmospheres 
means  a  pressure  of  forty-five  pounds  per  square  inch.  It 
is  this  pressure  which  causes  water  to  rise  in  pumps  and 
flow  through  siphons.  Atmospheric  pressure  is  exerted 
in  all  directions  and  is  variable.  It  is  measured  by  the 
barometer.  The  normal  or  standard  pressure  of  the  at- 
mosphere is  equal  to  the  weight  of  a  column  of  mercury 
one  square  inch  in  cross  section  and  29.92  in.  high,  or  one 
square  centimeter  in  cross  section  and  760  mm.  high.  But 
since  atmospheric  pressure  is  at  the  rate  of  fifteen  pounds  to 
the  square  inch,  it  is  necessary  to  know  the  height  only 
of  the  mercury  column  in  order  to  know  the  pressure. 

The  pressure  of  the  atmosphere  varies  as  the  height  and  the  compo- 
sition of  the  atmosphere  vary,  and  the  barometer  changes  accordingly. 
The  weight  of  a  liter  of  dry  air  at  o°  and  760  mm.  is  i  .293  gm. 

The  appreciable  movements  of  the  atmosphere  are  the  winds. 

Ingredients  of  the  Atmosphere.  —The  atmosphere  is  a 
mixture  of  several  gases.  But  since  this  mixture  always 
contains  about  78  parts  of  nitrogen  and  21  parts  of  oxygen 
by  volume,  we  often  speak  of  air  as  consisting  solely  of 
these  two  gases.  Besides  this  large  proportion  of  oxygen 
and  nitrogen,  the  air  always  contains  small  and  variable 
proportions  of  water  vapor  and  carbon  dioxide  gas.  Be- 
sides these  four  ingredients,  air  always  contains  the  gases 
argon  and  helium,  and  usually  ozone,  hydrogen,  hydrogen 
peroxide,  compounds  related  to  ammonia  and  nitric  acid, 
dust,  and  germs.  The  composition  varies  but  slightly  in 
different  localities.  Near  the  city  air  may  contain  a  rela- 
tively larger  proportion  of  dust,  ammonia,  sulphur  com- 
pounds, and  acids ;  in  the  country  the  proportion  of  ozone 
is  relatively  large ;  at  the  ocean  the  air  contains  consider- 
able salt. 


The  Atmosphere  —  Nitrogen.  63 

General  Properties  of  Nitrogen.  —  The  chemical  ele- 
ment, nitrogen,  constitutes  about  78  per  cent  of  the  atmos- 
phere (by  volume).  It  is  a  colorless  gas,  and  has  no  taste 
or  odor.  It  is  somewhat  lighter  than  air,  and  is  very 
slightly  soluble  in  water.  In  many  respects  it  differs 
markedly  from  oxygen.  Thus  it  will  not  support  combus- 
tion, neither  will  it  burn  nor  sustain  life.  Animals  die  if 

left  in  nitrogen. 

* 

The  fact  that  a  candle  flame  quickly  goes  out  and  a  mouse  soon  dies 
in  nitrogen  was  first  observed  by  Rutherford,  an  English  physician, 
who  discovered  the  gas  in  1772.  Soon  after,  Lavoisier  showed  the  true 
relation  of  nitrogen  to  the  atmosphere.  To  emphasize  the  inability 
of  the  gas  to  support  life,  he  called  the  new  gas  azote,  the  name  now 
used  for  it  by  some  French  chemists. 

Nitrogen  is  not  poisonous,  for  a  large  proportion  of  the 
air  we  breathe  is  nitrogen.  Its  function  in  the  atmosphere 
is  to  dilute  the  oxygen.  It  is  an  inert  element.  It  com- 
bines with  only  a  few  other  elements,  and  many  of  its 
compounds  easily  decompose. 

Oxygen  and  Nitrogen  in  the  Atmosphere.  —The  chem- 
ical activity  of  the  atmosphere  is  due  to  the  free  oxygen 
it  contains.  We  have  already  learned  that  oxygen  is  an 
i  active  chemical  element.  If  the  air  were  largely  oxygen, 
rusting  and  decay  would  proceed  with  astounding  rapidity, 
and  fires  once  started  would  burn  with  .great  violence.  On 
kthe  other  hand,  nitrogen  is  inactive.  And  if  the  air  con- 
tained much  more  than  the  normal  amount,  chemical 
action  would  be  slower.  Oxygen  alone  is  too  active, 
while  nitrogen  alone  is  inactive.  To  be  serviceable  to 
man,  oxygen  must  be  diluted  with  nitrogen,  while  nitro- 
gen must  be  accompanied  by  a  small  proportion  of 
oxygen. 


64  Descriptive  Chemistry. 

The  presence  of  oxygen  and  nitrogen  in  the  atmosphere,  and  the 
functions  of  the  two  gases,  were  first  clearly  explained  by  Lavoisier  in 
1 777,  though  many  others  —  Boyle,  Priestley,  Rutherford,  and  Scheele 
—  helped  solve  the  problem. 

Composition  of  the  Atmosphere.  —  Samples  of  air  from 
various  parts  of  the  globe  show  a  remarkable  uniformity 
of  composition.  Until  1895  it  was  supposed  that  pure  air 
consisted  solely  of  oxygen  and  nitrogen.  But  it  has  been 
found  that  about  one  per  cent  of  the  gas  hitherto  called 
nitrogen  is  argon,  a  gas  so  much  like  nitrogen,  and  so 
difficult  to  separate  from  the  latter,  that  for  years  it  had 
been  overlooked  (see  Argon,  below).  According  to  the 
most  recent  results,  the  following  is  — 

THE  COMPOSITION  OF  PURE  DRY  AIR. 


INGREDIENT. 

PERCENTAGE. 

By  volume. 

By  weight. 

Nitrogen 

78.06 
21.00 
0.94 

7#»Y 

-23.2 
•.V? 

Oxvefen 

Argon       .... 

•\ 

The  composition  of  the  atmosphere  was  studied  by  Priestley,  but  his 
results  were  conflicting.  Cavendish,  in  1781,  was  the  first  to  show  that 
the  proportion  of  oxygen  and  nitrogen  in  air  is  nearly  constant.  Since 
his  time  this  result  has  been  confirmed  by  many  chemists,  especially  by 
Bunsen,  who  is  widely  known  as  the  inventor  of  the  Bunsen  burner, 
which  is  used  as  a  source  of  heat  in  chemical  laboratories. 

The  Volumetric  Composition  of  the  Air  may  be  found 
by  introducing  a  known  volume  of  pure  air  into  a  eudiom- 
eter and  exploding  it  with  a  known  volume  of  hydrogen. 
The  oxygen  of  the  air  combines  with  twice  its  volume  of 
hydrogen,  forming  a  minute  quantity  of  water ;  hence  one 


The  Atmosphere  —  Nitrogen. 


third  of  the  diminution  in  volume  is  the  volume  of  oxygen 
in  the  air.  The  difference  between  the  volume  of  oxygen 
found  and  the  original  volume  of  air  is  the  volume  of 
nitrogen. 

An  illustration  will  make  this  experiment  clear.  Suppose  (i)  we 
mix  and  explode  loocc.  of  air  and  50  cc.  of  hydrogen,  or  15000.  in  all, 
and  (2)  that  the  residue  measures  87  cc.  Now,  150  —  87  =  63,  hence 
63  cc.  of  the  total  volume  combined  to  form  water.  But  one  third  of 
63  cc.  is  oxygen,  which  came  from  the  original  volume  of  air.  Hence, 
63  -r-  3  =  21,  the  volume  of  oxygen  in  100  cc.  of  air.  The  remainder, 
79  cc.,  is  nitrogen,  argon, 
and  other  gases. 

Another  Method,  <;often 
used  to  determine  the  volu- 
metric composition  of  the 
air,  is  based  on  the  fact 
that  phosphorus  will  com- 
bine slowly  with  oxygen, 
even  at  the  ordinary  tem- 
perature. The  operation  is 
performed  in  an  apparatus 
like  that  shown  in  Figure  1 3. 
A  piece  of  phosphorus,  C, 
attached  to  a  wire,  is 
inserted  into  a  graduated 
glass  tube,  />,  containing  a 


measured  volume  of  air. 
White  fumes  indicate  im- 
mediate action.  These 
fumes  are  solid  particles 
of  an  oxide  of  phosphorus 


FIG.  13.  —  Apparatus  for  determining  the  com- 
position of  air  by  phosphorus. 


called  phosphorus  pentoxide.  'They  soon  dissolve  in  the  water,  which 
rises  higher  in  the  tube,  as  the  oxygen  combines  with  the  phosphorus. 
In  a  few  hours  the  phosphorus  is  removed,  and  the  volume  of  gas  is 
read.  The  difference  between  the  first  and  last  volumes  is  oxygen.  The 
gas  remaining  in  the  tube  is,  of  course,  a  mixture  of  nitrogen  and  argon. 
In  performing  this  experiment  unusual  care  must  be  taken  not  to  touch 
the  phosphorus  with  the  bare  hands. 


66  Descriptive  Chemistry. 

The  Gravimetric  Composition  of  Air  was  first  accurately 
determined  in  1841  by  the  French- chemists,  Dumas  and 
Boussingault.  The  average  result  of  many  experiments 

•tTT'O  o    Lm 

Oxygen   .     .     .     .     23  parts  by  weight. 
Nitrogen      ...     77  parts  by  weight. 

We  know,  however,  that  the  correct  proportions  are  — 

Oxygen  ....  23.2  parts  by  weight. 
Nitrogen  .  .  .  75.5  parts  by  weight. 
Argon  ....  1.3  parts  by  weight. 

They  passed  pure  air  through  a  weighed  tube  containing  copper,  and 
arranged  so  that  heat  could  be  applied.  The  oxygen  of  the  air  com- 
bined with  the  copper,  while  the  nitrogen  passed  on  into  a  weighed  globe. 
Both  tube  and  globe  increased  in  weight.  The  increase  in  the  tube  was 
the  weight  of  the  oxygen,  while  the  increase  in  the  globe  was  the  weight 
of  the  nitrogen. 

Water  Vapor  in  the  Atmosphere.  —  Water  vapor  is 
always  present  in  the  atmosphere,  owing  to  the  constant 
evaporation  from  the  ocean  and  other  bodies  of  water. 
The  total  amount  present  is  large,  though  variable.  A 
given  volume  of  air  will  absorb  a  definite  volume  of  water 
vapor  and  no  more,  and  the  amount  depends  largely  upon 
the  temperature.  Air  containing  its  maximum  amount  of 
water  vapor  is  said  to  be  saturated  at  that  temperature,  or 
to  contain  100  per  cent  of  water  vapor.  The  saturation 
point  is  also  called  the  dew  point.  On  a  pleasant  day  the 
relative  humidity  of  the  air,  i.e.  the  amount  of  water 
vapor  present,  may  vary  from  30  to  90  per  cent,  the  aver- 
age being  about  50  per  cent.  Warm  air  holds  more  vapor 
than  cool  air.  The  amount  of  water  vapor  in  the  air  has 
a  marked  influence  on  the  physical  condition  of  man. 
The  depressing  weather  during  "  dog  days  "  is  due  to  the 


The  Atmosphere  —  Nitrogen.  67 

high  relative  humidity  of  the  air,  which  sometimes  reaches 
95  per  cent.  The  absence  of  life  in  deserts  is  largely  due  * 
to  the  dry  air  .above  them.  Much  of  the  languor  felt  in  a 
"  close "  room  or  crowded  hall  is  partly  caused  by  the 
excess  of  water  vapor  in  the  "bad"  air.  The  presence  of 
water  vapor  in  the  air  is  shown  by  the  moisture  which  col- 
lects on  the  outside  of  a  vessel  containing  cold  water,  such 
as  a  pitcher  of  iced  water.  The  moisture  comes  from  the 
air  around  the  vessel.  For  a  similar  reason,  water  pipes 
in  a  cellar  and  the  cellar  walls  themselves  are  moist  in 
summer.  The  deliquescence  of  calcium  chloride,  common 
salt,  and  other  substances  likewise  reveals  the  presence  of 
water  vapor  in  the  air  (see  Deliquescence). 

When  the  temperature  of  the  air  falls,  the  water  vapor  condenses  and 
is  deposited  in  the  form  of  dew,  rain,  fog,  mist,  frost,  snow,  sleet,  or 
hail.  The  clouds  are  masses  of  water  vapor  which  has  been  condensed 
by  the  cold  upper  air. 

Carbon  Dioxide  in  the  Atmosphere.  —  Carbon  dioxide 
is  one  product  of  the  respiration  of  animals,  and  of  the 
combustion  and  decay  of  organic  substances.  By  these 
processes  an  immense  quantity  of  carbon  dioxide  is  being 
constantly  poured  into  the  atmosphere.  The  quantity  in 
the  atmosphere  is  variable,  though  not  between  such  wide 
limits  as  the  water  vapor.  The  proportion  in  normal  air 
is  about  4  parts  in  10,000  parts  of  air.  Over  the  ocean 
the  proportion  is  smaller,  but  in  the  air  of  cities  it  is 
greater.  In  crowded  rooms  the  proportion  is  often  as 
high  as  33  parts  in  10,000,  because  carbon  dioxide  is 
exhaled  faster  than  it  can  be  removed.  The  proportion  of 
carbon  dioxide  in  the  atmosphere  as  a  whole  is  practically 
constant,  largely  owing  to  the  fact  that  this  gas  is  an 
essential  food  of  plants  (see  Carbon  Dioxide).  The  pres- 
ence of  carbon  dioxide  in  the  air  is  detected  by  limewater. 


68  Descriptive  Chemistry. 

If  Hmewater  is  exposed  to  the  air,  the  carbon  dioxide  unites  with  the 
lime  in  the  limewater,  forming  a  thin,  white  crust  of  insoluble  calcium 
carbonate  on  the  surface  of  the  Hmewater.  If  air  is  drawn  through  lime- 
water,  the  liquid  becomes  milky,  because  the  particles  of  calcium  carbon- 
ate are  suspended  in  the  liquid.  The  purity  of  air  is  often  determined 
by  finding  out  what  proportion  of  carbon  dff>xide  it  contains.  If  a 
known  volume  of  dry  air  is  drawn  through  a  known  weight  of  Hmewater 
or  similar  liquid,  the  increase  in  weight  will  be  the  weight  of  carbon 
dioxide  in  the  volume  of  air  used. 

The  different  gases  in  the  atmosphere  are  not  arranged 
in  layers  according  to  their  densities.  They  are  in  con- 
stant circulation  (see  Diffusion).  Hence  carbon  dioxide, 
though  heavier  than  oxygen  and  nitrogen  (volume  for  vol- 
ume), does  not  remain  nearest  the  ground,  but  is  distrib- 
uted through  the  air.  In  a  few  exceptional  localities, 
carbon  dioxide  arises  from  volcanoes  faster  than  it  can 
diffuse,  and  fills  the  adjacent  valley. 

Argon  in  the  Atmosphere.  —  Argon  is  a  colorless,  odor- 
less gas.  Its  chief  characteristic  is  its  chemical  inactivity. 
No  compounds  of  argon  have  as  yet  been  prepared  or 
discovered.  The  name  argon  is  happily  chosen,  being 
derived  from  Greek  words  signifying  inert.  It  constitutes 
0.94  per  cent  by  volume  of  the  atmosphere,  or  1.3  per 
cent  by  weight. 

Argon  was  discovered  in  1894  by  Rayleigh  and  Ramsay.  Rayleigh 
had  found  that  nitrogen  from  air  weighed  more  than  an  equal  volume 
of  nitrogen  obtained  from  compounds  of  nitrogen.  Consequently,  they 
believed  that  the  nitrogen  from  air  contained  another  gas  hitherto  over- 
looked. A  series  of  elaborate  experiments  showed  that  after  all  the 
oxygen  and  nitrogen  was  removed  from  purified  air,  there  still  remained 
a  small  quantity  of  a  new  gas,  which  they  called  argon.  It  may  be  pre- 
pared (i)  by  passing  pure  air  over  healed  copper  to  remove  the  oxygen, 
and  then  the  remaining  gas  over  heated  magnesium  or  calcium  to  remove 
the  nitrogen ;  or  (2)  by  passing  electric  sparks  through  a  mixture  of  air 
and  oxygen,  and  removing  the  compound  of  oxygen  and  nitrogen  as  fast 


The  Atmosphere  —  Nitrogen.  69 

as  it  is  formed.  The  latter  method  is  a  repetition  of  the  one  used  by 
Cavendish  when  he  determined  the  composition  of  air,  and  he  would 
have  no  doubt  discovered  argon  had  he  continued  his  investigations. 

Inert  Gases  in  the  Atmosphere.  —  Helium,  neon,  krypton,  and  xenon 
have  recently  been  discovered  by  Ramsay.  At  present  little  is  known 
about  these  gases.  They  resemble  argon  in  being  inactive  chemical 
elements.  They  constitute  an  exceedingly  minute  proportion  of  the 
atmosphere.  Helium  is  also  found  in  certain  rare  minerals,  in  the  gases 
from  some  mineral  springs,  and  in  the  atmosphere  of  the  sun.  It  is 
about  twice  as  heavy  as  hydrogen.  According  to  Ramsay,  "  it  is  prob- 
able that  helium  is  continually  escaping  from  the  earth  in  small  quantities 
in  certain  regions.'1 

Air  is  a  Mixture,  in  spite  of  the  fact  that  we  speak  of 
its  "composition."  Chemical  compounds  have  two  invari- 
able characteristics  :  viz.,  (i)  their  components  are  in  a  fixed 
proportion,  and  (2)  their  formation  and  decomposition  are 
usually  attended  by  definite  evidences  of  chemical  action, 
such  as  light,  heat,  change  of  color  and  form,  etc.  The 
following  facts  show  that  air  is  a  mixture  of  free  gases :  — 

(1)  The  proportion  of   oxygen   and   of  nitrogen   is  not 
fixed,    but   varies   between    small    limits,    which    may   be 
detected  by  accurate  analysis. 

(2)  When  nitrogen  and  oxygen  are  mixed  in  the  propor- 
tions which  form  air,  the  product  is  exactly  like  air,  but 
the  act  of  mixing  gives  no  evidence  of  chemical  action. 

(3)  When  air  is  dissolved  in  water,  a  greater  proportion 
of  oxygen  than  of  nitrogen  dissolves.     If  the  oxygen  and 
nitrogen  were  combined  in  the  air,  the  dissolved  air  would, 
of  course,  have  the  same  composition  as  air  itself. 

Liquid  Air  is  a  mixture  of  the  liquefied  gases  which  con- 
stituted the  air  used.  It  is  a  milky  liquid,  owing  to  the 
presence  of  solid  carbon  dioxide  and  ice.  If  these  solids 
are  removed  by  filtering,  the  filtrate  has  a  pale  blue  tint. 
It  is  slightly  heavier  than  water.  It  is  intensely  cold,  its 


jo  Descriptive  Chemistry. 

temperature  being  about  —200°  C.  It  boils  at  about 
— 190°  C.  under  atmospheric  pressure.  If  a  tumbler  is 
filled  with  liquid  air,  the  latter  boils  vigorously,  the  sur- 
rounding air  becomes  intensely  cold,  frost  gathers  on  the 
tumbler,  and  in  a  short  time  the  liquid  air  will  have 

entirely  disappeared  into 
the  air  of  the  room.  If, 
however,  the  liquid  air  is 
placed  in  a  Dewar's  bulb 
or  flask,  it  evaporates  so 
slowly  that  some  will  remain 
in  the  flask  several  hours. 

The  Dewar's  bulb  (Fig.  14) 
consists  of  two  flasks,  one  within 
the  other,  attached  at  the  top  ;  the 
space  between  the  flasks  is  a 
vacuum.  Sometimes  the  outer 
surface  of  the  inner  flask  is  coated 
with  mercury  or  silver,  which 
helps  to  protect  the  liquid  air  from 
the  heat  of  the  atmosphere.  In 
transporting  liquid  air  a  large 
Dewar's  bulb  or  similar  device  is 
FIG.  14. — A  Dewar's  bulb.  used.  One  form  consists  of  a 

large    metal    can    wrapped    with 

many  thicknesses  of  felt  and  slipped  into  a  larger  can  covered  with 
canvas  or  felt.  The  liquid  air  is  put  in  the  inner  can  and  a  loose 
stopper  or  piece  of  felt  is  placed  over  the  mouth.  The  liquid  may  also 
be  kept  in  these  cans  for  some  time  with  only  a  moderate  loss,  unless 
the  surrounding  temperature  is  exceptionally  high. 

Liquid  air,  owing  to  its  extremely  low  temperature,  pro- 
duces remarkable  physical  changes.  A  tin  or  iron  vessel 
which  has  been  cooled  by  liquid  air  is  so  brittle  that  it  may 
often  be  crushed  with  the  fingers.  Nearly  all  plastic  or 
soft  substances,  including  many  kinds  of  food,  when  im- 


The  Atmosphere  —  Nitrogen.  71 

mersed  in  liquid  air,  become  hard  and  brittle,  leather  being 
the  only  important  exception.  Mercury  freezes  so  hard  in 
liquid  air,  that  it  may  be  used  as  a  hammer  to  drive  a  nail. 
When  liquid  air  is  put  in  a  teakettle  standing  on  a  block 
of  ice,  the  liquid  air  boils  vigorously.  If  the  kettle  of 
liquid  air  is  placed  over  a  lighted  Bunsen  burner,  frost  and 
ice  collect  on  the  bottom  of  the  kettle,  because  the  intense 
cold  of  the  kettle  solidifies  the  water  vapor  and  carbon 
dioxide,  which  are  the  two  main  products  of  burning 
illuminating  gas.  If  water  is  now  poured  into  the  kettle, 
the  liquid  air  boils  over  and  the  water  is  instantly  frozen ; 
the  water  is  so  much  hotter  than  the  liquid  air  that  the  latter 
boils  more  violently,  and  since  its  rapid  evaporation  causes 
absorption  of  heat,  the  water  gives  up  its  heat  and  becomes 
ice.  Ordinary  liquid  air*  is  from  one  half  to  one  fifth  liquid 
oxygen,  and  will  support  combustion.  A  red-hot  rod  of 
steel  or  of  carbon  burns  brilliantly  in  this  cold  liquid. 

Numerous  applications  of  liquid  air  have  been  proposed,  but  thus  far 
they  have  not  passed  the  experimental  stage.  It  has  been  suggested 
that  it  be  used  as  a  refrigerant  instead  of  ice,  for  ventilating  and  cooling 
rooms,  as  a  blasting  material,  for  removing  diseased  flesh  from  a  wound, 
for  destroying  refuse,  and  as  a  commercial  source  of  oxygen.  The  last 
use  is  based  primarily  on  the  fact  that  as  liquid  air  evaporates,  the 
nitrogen  passes  off  first,  and  in  a  short  time  relatively  pure  oxygen 
remains  (see  Oxygen). 

A  little  liquid  air  was  produced  in  1883  with  considerable  labor  and 
at  an  enormous  expense.  Now  it  is  Easily  manufactured  in  large  quan- 
tities at  a  comparatively  low  cost.  In  the  older  methods  of  preparing 
liquefied  gases,  the  gas  was  subjected  to  tremendous  pressure  and  a  low 
temperature.  At  present,  air  is  liquefied  by  a  different  method.  Com- 
pressed air  cooled  by  water  is  forced  through  a  pipe  with  a  small  open- 
ing into  a  larger  cylinder  called  the  liquefier.  As  it  escapes  into  the 
liquefier  it  expands  and  its  temperature  falls,  because  expansion  is  a 
cooling  process.  The  temperature  of  the  liquefier  is  thus  reduced,  so 
that  the  air,  which  continues  to  enter,  expands  at  such  a  low  temperature 
that  it  becomes  a  liquid. 


72  Descriptive  Chemistry. 

NITROGEN. 

Occurrence.  —  Nitrogen,  besides  comprising  four  fifths 
of  the  atmosphere,  is  a  component  of  nitric  acid  and  am- 
monia, and  of  the  many  compounds  related  to  them.  It 
is  also  an  essential  constituent  of  animal  and  vegetable 
matter. 

The  name  nitrogen  was  given  to  the  gas  by  Chaptal  from  the  fact 
that  it  is  a  component  of  niter,  an  old  name  of  potassium  nitrate. 

Preparation.  —  Nitrogen  is  usually  obtained  from  the  air  by  remov- 
ing the  oxygen  by  phosphorus.  A  tall  jar  is  placed  over  burning 
phosphorus  contained  in  a  shallow  dish  floating  in  a  large  vessel  of 
water.  The  oxygen  combines  with  the  phosphorus,  leaving  nitrogen, 
more  or  less  pure,  in  the  jar.  Other  methods  may  be  used,  such  as 
decomposing  ammonium  nitrite  by  heat,  or  passing  air  over  heated 
copper. 

Additional  Properties.  —  In  addition  to  its  inertness,  already  men- 
tioned, nitrogen  is  a  little  lighter  than  air,  and  is  very  sparingly  soluble 
in  water.  Its  density  is  0.972  (air  =  i).  One  liter  at  o°  C.  an'd  760  mm. 
weighs  i.256gm.  One  hundred  liters  of  water  dissolve  only  1.5  1.  at  the 
ordinary  temperature.  It  combines  with  magnesium  and  a  few  other 
metals  at  a  red  heat,  forming  nitrides.  Electric  sparks  cause  nitrogen  to 
combine  with  oxygen  and  with  hydrogen,  forming  ultimately  nitric  acid 
and  ammonia,  hence  these  substances  or  others  related  to  them  are 
often  found  in  the  rain  which  falls  during  a  thunder  storm. 

Relation  of  Nitrogen  to  Life.  —  Oxygen,  carbon  diox- 
ide, and  water  vapor  are  essentially  related  to  the  life  of 
plants  and  animals.  Nitrogen  is  also  vitally  connected 
with  different  forms  of  life.  Atmospheric  nitrogen  merely 
dilutes  the  oxygen.  Although  we  live  in  an  atmosphere 
containing  such  a  large  proportion  of  nitrogen,  we  cannot 
assimilate  it.  According  to  a  reliable  authority,  "  the  air 
as  it  leaves  the  lungs  contains  79.5  per  cent  of  nitrogen," 
and  hence  cannot  become  a  part  of  the  body.  Yet  all  flesh 
contains  nitrogen,  and  the  rejected  waste  products  of  ani- 


The  Atmosphere  —  Nitrogen.  73 

mals  are  largely  combined  nitrogen.  The  nitrogen  needed 
by  animals  must  be  in  combination  to  become  available. 
And  it  is  taken  in  the  form  of  nitrogenous  food,  such  as 
lean  meat,  fish,  wheat  and  other  grains. 

Most  plants  take  up  combined  nitrogen  from  the  soil  in 
the  form  of  nitrates  (compounds  derived  from  nitric  acid) 
or  of  ammonia.  Hence  combined  nitrogen  is  being  con- 
stantly taken  from  the  soil,  and  in  order  to  preserve  the 
fertility  of  the  soil,  nitrogen  must  be  supplied.  This  is  done 
by  allowing  nitrogenous  organic  matter  to  decay  upon  the 
soil,  or  by  adding  to  the  soil  a  fertilizer,  which  is  a 
mixture  containing  nitrogen  compounds.  Recently  it 
has  been  shown  that  leguminous  plants,  such  as  peas, 
beans,  and  clover,  take  up  nitrogen  from  the  air  by  means 
of  bacteria,  which  are  in  nodules  on  their  roots. 

EXERCISES. 

1.  What  is  the  atmosphere?     What  is  air?     What  is  the  literal 
meaning  of  the  word  atmosphere?    What  is  the  wind? 

2.  Develop  the  topics:    (a)  atmospheric  pressure,  (b)  occurrence  of 
nitrogen,  (c)  volumetric  composition  of  the  air,  (W)  gravimetric  com- 
position of  the  air,   (e)  water  vapor  in  the  atmosphere,    (/")  carbon 
dioxide  in  the  atmosphere,  (g)  air  is  a  mixture. 

3.  Define  and  illustrate  the  terms  :    (#)  an  atmosphere,  (<£)  normal 
pressure,   (c)  standard  pressure,   (d)  dew  point,  (e)  relative  humidity, 
(/)  inert. 

4.  What  are  the  two  chief  ingredients  of  the  atmosphere?     The  per- 
manent ingredients  ?     The  variable  ingredients  ?    The  ingredients  found 
in  traces?     What  are  sometimes  found  in  the  air  of  cities? 

5.  What  is  the  symbol  of  nitrogen?     What  are  its  general  proper- 
ties?    Its  special  properties?     What  is  its  main  function  in  the  atmos- 
phere?    How  may  it  be  prepared? 

6.  When  and  by  whom  was  nitrogen  discovered?     Why  and  by 
whom  was  it  named  "azote11  and  "nitrogen11  ? 

7.  What  is  the  relation  of  nitrogen  to  animal  and  to  vegetable  life? 


74  Descriptive  Chemistry. 

8.  Compare  the  functions  of  oxygen  and  nitrogen  in  the   atmos- 
phere.    What  famous  chemists  helped  solve  this  problem? 

9.  State  the  composition  of  pure  air  (a)  by  volume,  and  (b)   by 
weight. 

10.  Give  a  brief  biographical  account  of  (a)  Cavendish,  (^)  Dumas, 
(c)  Rutherford.     (See  Appendix,  §  4.) 

11.  What  is  a  cloud?     The  dew?     Why  does  moisture  gather  on 
cellar  walls?     Why  are  mines  often  damp?     What  is  (a)  rain,  (t>)  fog, 

(c)  mist? 

12.  Describe  the  action  of  air  upon  (a)  limewater,  and  (b}  calcium 
chloride. 

13.  How  does  the  atmosphere  illustrate  the  diffusion  of  gases? 

14.  What  is  argon?     Give  a  brief  account  of  (a)  its  discovery,  (£)  its 
properties,  (c)  its  method  of  preparation.     What  proportion  of  pure 
air  is  argon?     What  is  the  significance  of  the  name  argon  f 

15.  Give  a  brief  account  of  helium,  neon,  krypton,  and  xenon. 

1 6.  What   is   liquid   air?     What   are   its   chief  properties?     State 
briefly  its  method  of  manufacture.     Describe  its  action  (a)  upon  solids, 
such  as  rubber,  (b)  upon  liquids,  such  as  mercury,  (c)  upon  hot  steel, 

(d)  when  evaporated  quickly.     Describe  a  Dewar's  bulb. 


PROBLEMS. 

1.  If  a  man  inhales  18  cu.  ft.  of  air  an  hour,  what  weight  of  oxy- 
gen does  he  consume  in  24  hr.  ? 

2.  What  is  the  weight  of  air  in  a  room,  6x6x3111.,  if  a  liter  of 
the  air  weighs  1.3  gm.  ? 

3.  A  mixture  of  25  cc.  of  air  and  50  cc.  of  hydrogen  is  exploded. 
The  residue  measures  60.3  cc.     What  per   cent  of  oxygen   did   this 
sample  of  air  contain  ? 

4.  How  many  kilograms  of  pure  air  are  needed  to  yield  100  kg.  of 
oxygen  ? 

5.  Express  in  inches  the  following  barometer  readings :   (a)  760 
mm.,  (<£)  740  mm.,  (c)  75  cm.,  (d)  0.749  m.,  (e)  7.67  dm. 

6.  Dumas  and  Boussingault,  in  1841,  found  in  a  sample  of  air, 
12 -373  gm-  of  nitrogen  and  3.68  gm.  of  oxygen.     What  per  cent  of 
each  was  found? 

7.  What  is  the  weight  at  o°  C.  and  760  mm.  of  (#)  1000  cc.  of  dry 
air?     Of  (£)  750  1.,  (c)  1750  cc.,  (d)  850  cu.  m.? 


CHAPTER   VII. 

LAW  AND  THEORY  — LAWS  OF  DEFINITE  AND  MUL- 
TIPLE PROPORTIONS  —  ATOMIC  THEORY  — ATOMS  AND 
MOLECULES  — SYMBOLS  AND  FORMULAS  — EQUATIONS. 

Law  and  Theory.  —  We  discover  facts  by  observation 
and  experiment.  Facts  which  always  oc"cur  under  the 
same  circumstances  soon  become  well  established.  Such 
facts  are  often  -summarized  in  a  brief  statement  called  a 
law. 

Sometimes  the  word  law  is  used  in  the  sense  of  the  uniform  behavior 
summarized  in  the  brief  statement.  Hence,  in  a  narrow  sense,  a  law 
is  a  statement  of  a  fact,  but  in  a  broad  sense  a  law  is  the  fact  itself. 
Thus,  the  law  of  definite  proportions  (soon  to  be  discussed)  is  either 
(i)  a  brief  statement  of  the  general  fact  of  definite  proportions  of  ele- 
ments in  compounds,  or  (2)  the  uniform  behavior  itself  as  far  as  the 
composition  of  chemical  compounds  is  concerned. 

The  cause  of  many  scientific  facts  is  unknown.  The 
explanation  we  give,  or  the  statement  we  make,  of  the 
cause  of  facts  is  called  a  theory.  Laws  are  statements  of 
fact,  theories  are  statements  of  the  supposed  cause  of  facts. 
Thus  we  know  that  chemical  compounds  have  a  definite 
composition,  because  we  have  discovered  by  experiment 
the  facts  on  which  this  law  is  based ;  and  we  have  framed 
a  theory,  which,  as  far  as  our  present  knowledge  is.  con- 
cerned, is  a  satisfactory  explanation  of  the  cause  of  the 
general  fact  of  definite  composition.  Laws  seldom  change, 
but  theories  are  often  modified.  Laws  are  the  result  of 
experiment,  theories  are  the  outcome  of  mental  operations. 

75 


76  Descriptive  Chemistry. 

We  accept  a  certain  theory  until  a  more  satisfactory  one  is 
proposed.  If  a  fact  is  not  well  established  or  is  not  gen- 
eral, we  account  for  it  by  an  hypothesis.  An  hypothesis 
is  a  guess  or  supposition  concerning  the  cause  of  some 
particular  fact  or  set  of  facts,  and  it  is  usually  proposed  as 
a  basis  for  making  further  experiments.  Hypotheses  often 
lead  to  theories. 

Laws,  theories,  and  hypotheses  are  of  great  service  in 
chemistry,  since  they  help  us  gather  into  intelligible  state- 
ments a  vast  number  of  facts  which  are  apparently  not 
related.  They  also  assist  in  discovering  facts. 

Law  of  Definite  Proportions  by  Weight.  —  When  the 
metal  magnesium  is  heated  in  the  air,  it  burns  with  a 
dazzling  flame  into  a  grayish  powder,  due  to  combination 
with  oxygen.  If  a  known  weight  of  magnesium  is  heated 
in  a  crucible,  so  that  the  product  cannot  escape,  a  remark- 
able relation  is  revealed.  In  order  to  burn  completely  1.5 
gm.  of  magnesium,  i  gm.  of  oxygen  is  necessary;  and 
the  product,  magnesium  oxide,  weighs  2.5  gm.  This 
product  contains,  therefore,  60  per  cent  magnesium  and 
40  per  cent  oxygen.  Accurate  repetitions  of  this  experi- 
ment have  shown  that  this  proportion  by  weight  is  fixed 
and  definite.  Again,  if  all  the  oxygen  is  driven  from  a 
weighed  quantity  of  potassium  chlorate  by  heating  this 
compound  in  a  crucible,  39.18  per  cent  of  oxygen  is 
always  obtained.  This  means  that  the  proportion  of 
potassium,  chlorine,  and  oxygen  which  makes  up  potas- 
sium chlorate  is  fixed  and  definite.  Otherwise,  the  prop- 
erties of  potassium  chlorate  would  vary.  Experiments 
similar  to  these  show  that  in  all  chemical  compounds  the 
different  components  are  always  present  in  a  definite  and 
unvarying  proportion  by  weight.  There  are  no  exceptions 
to  this  general  fact.  This  constancy  of  proportion  in 


Law  of  Multiple  Proportions.  77 

chemical  compounds  is  stated  as  the  Law  of  Definite  Pro- 
portions by  Weight,  thus :  - 

A  given  chemical  compound  always  contains  the  same 
elements  in  the  same  proportions  by  weight. 

Sometimes  it  is  condensed  into  this  form :  — 

A  chemical  compound  has  a  definite  composition  by  weight. 

This  law  is  one  of  the  fundamental  laws  of  chemistry.  It  is  so  firmly 
believed  that  if  the  composition  of  a  compound  is  found  by  analysis  to 
vary,  chemists  conclude  that  the  experimental  work  is  incorrect  or  that 
the  compound  is  impure.  The  law  was  established  as  the  outcome  of 
a  controversy  between  two  French  chemists,  Proust  (1755-1826)  and 
Berthollet  (1748-1822).  The  discussion  lasted  from  1799  to  1806. 
Berthollet  believed  that  compounds  might  have  a  varying  composition. 
Indeed,  by  his  experiments  he  detected  "  gradual  changes "  in  com- 
position. But  Proust  showed  that  Berthollet  analyzed  mixtures  and 
not  compounds.  In  a  mixture  the  parts  may  be  present  in  any  propor- 
tion. Subsequent  experiments  have  only  strengthened  our  confidence 
in  this  law. 

Law  of  Multiple  Proportions.  —  Proust  showed  that 
some  elements  combine  in  more  than  one  proportion, 
and  thereby  produce  distinct  compounds.  But  he  failed  to 
notice  that  if  the  weight  of  one  element  is  constant,  the 
varying  weights  of  the  other  element  are  in  a  simple  mul- 
tiple relation  to  each  other.  Dalton  discovered  this  gen- 
eral fact  about  1804.  The  composition  of  compounds  is 
usually  expressed  in  per  cent ;  but  such  expressions  in  a 
series  of  compounds  reveal  nothing  about  multiple  rela- 
tions. If,  however,  a  constant  weight  is  adopted  as  a  unit 
for  one  component,  and  the  composition  of  the  series  of 
compounds  is  expressed  in  terms  of  this  unit,  then  the 
simple  multiple  relation  which  exists  between  the  weights 
of  the  other  component  is  clearly  seen.  Thus,  we  learn> 
little  from  the  statement  that  the  two  compounds  of  carbon 


Descriptive  Chemistry. 


and  oxygen  contain  73  and  57  per  cent  of  oxygen.  But 
if  in  expressing  the  composition  of  these  compounds 
we  adopt  12  as  the  weight  of  carbon,  the  weights  of 
oxygen  become  32  and  16,  i.e.  the  weights  of  oxygen  are 
simple  multiples.  The  five  compounds  of  oxygen  and 
nitrogen,  which  will  soon  be  studied,  aptly  illustrate  this 
fact : — 

TABLE 'TO  ILLUSTRATE  MULTIPLE  PROPORTIONS. 


COMPOSITION  IN 

UNIT 

PER  CENT. 

WEIGHT. 

RATIO. 

NAME. 

Nitrogen.     Oxygen. 

Nitrogen. 

Nitroj 

;en.     Oxygen. 

Nitrous  oxide    .... 

63.6            36,4 

7 

7 

4 

Nitric  oxide 

j.6  6        zi  A. 

7 

7 

8 

Nitrogen  trioxide  .     .     . 

36.8        63.2 

7 

7 

12 

Nitrogen  peroxide.      .     . 

30.4        69.6 

7 

7 

16 

Nitrogen  pentoxide    . 

25.9        74.1 

7 

7 

20 

From  this  table  it  is  clear  that  the  weights  of  oxygen 
combined  with  the  same  weight  of  nitrogen  are  as  1:2: 
3:4:5,  i.e.  they  are  simple  multiples  of  each  other. 

The  general  fact  of  multiple  proportions  is  expressed  in 
the  Law  of  Multiple  Proportions,  thus :  - 

When  two  or  more  elements  unite  to  form  a  series  of 
compounds,  a  fixed  weigJit  of  one  element  so  combines  with 
different  weights  of  the  other  element  that  the  relations  be- 
tween the  different  weights  can  be  expressed  by  small  whole 
numbers. 

This  law,  like  the  law  of  definite  proportions,  is  a  fun- 
damental law  of  chemistry,  and  together  they  have  pro- 
foundly influenced  its  theoretical  and  practical  progress. 


JOHN    DALTON 

1766-1844 

THE    ENGLISH    CHEMIST   WHO    LAID    THE    FOUNDATIONS   OF    THEORETICAL    CHEMISTRY 


The  Atomic  Theory.  79 

The  Atomic  Theory  of  the  constitution  of  matter  was 
proposed  by  Dalton  to  explain  the  laws  of  definite  and 
multiple  proportions.  This  theory  assumes  (i)  that  the 
chemical  elements  consist  ultimately  of  a  vast  number  of 
very  small,  indivisible  particles  or  atoms,  (2)  that  the 
atoms  of  the  same  element  have  the  same  weight,  (3)  that 
atoms  of  different  elements  have  different  weights,  and  (4) 
that  chemical  action  is  union  or  separation  of  the  atoms  of 
the  elements. 

Let  us  now  consider  how  this  theory  explains  the  facts 
summarized  in  the  laws  of  definite  and  multiple  propor- 
tions, (i)  When  magnesium  combines  with  oxygen,  1.5 
parts  by  weight  of  magnesium  combine  with  one  part  by 
weight  of  oxygen.  Analysis  of  the  product  —  magnesium 
oxide  —  shows  that  this  proportion  is  constant;  that  is, 
•pure  magnesium  oxide  always  contains  the  elements  mag- 
nesium and  oxygen  in  this  proportion.  Now,  according  to 
the  atomic  theory,  magnesium  oxide  is  the  product  of  the 
union  of  indivisible  atoms  of  magnesium  and  indivisible 
atoms  of  oxygen.  It  therefore  follows  that  when  magne- 
sium and  oxygen  unite,  atom  for  atom,  the  magnesium 
oxide  must  contain  the  two  elements  in  the  proportion  of 
the  weights  of  their  atoms,  i.e.  it  must  always  have  the 
same  composition.  It  is  immaterial  whether  the  actual 
weights  of  these  elements  which  combine  are  in  the  pro- 
portion of  i  to  1.5,  because  whatever  is  in  excess  of  this 
proportion  will  be  left  uncombined.  For  example,  if  we 
start  with  i  gm.  of  oxygen  and  2  gm.  of  magnesium,  then 
0.5  gm.  of  magnesium  will  be  left  uncombined.  Thus  the 
atomic  theory  explains  the  law  of  definite  proportions.  (2) 
But  atoms  do  not  always  combine  in  the  simple  proportion 
of  i  to  i.  They  may  combine  in  the  proportions  of  i  to  2, 
2  to  3,  i  to  3,  i  to  4,  etc.  But  according  to  the  atomic 


8o  Descriptive  Chemistry. 

theory  atoms  are  assumed  to  be  indivisible.  Hence,  if  we 
assume  the  atomic  theory,  the  proportions  of  the  weights 
of  different  elements  in  a  series  of  compounds  must  be 
simple  proportions,  i.e.  the  elements  must  unite  in  accord- 
ance with  the  law  of  multiple  proportions.  To  illustrate : 
There  are  two  compounds  of  carbon  and  oxygen.  Since 
atoms  are  indivisible,  the  simplest  combinations  of  the  atoms 
are  (i)  one  atom  of  carbon  to  one  atom  of  oxygen,  and  (2) 
one  atom  of  carbon  to  two  atoms  of  oxygen.  Analysis 
shows  that  in  the  first  compound  the  proportion  of  carbon 
to  oxygen  is  6  to  8.  According  to  the  theory,  the  propor- 
tion in  the  second  compound  should  be  6  to  16;  this  pro- 
portion is  verified  by  analysis.  In  other  words,  if  we 
adopt  6  as  the  weight  of  carbon  in  its  two  oxides,  then  the 
weights  of  oxygen  are  in  the  simple  proportion  i  to  2. 

Atoms  and  Molecules.  —  It  should  not  be  forgotten  that 
the  laws  of  definite  and  multiple  proportions  deal  with 
facts,  and  that  the  atomic  theory  deals  with  conceptions 
which  may  be  true,  but  which  cannot  be  proved  to  be 
true.  We  often  speak  of  atoms  as  if  they  could  be  per- 
ceived by  the  senses,  but  we  do  so  simply  because  such 
expressions  help  us  describe,  study,  and  interpret  chemical 
action.  According  to  the  present  views,  atoms  do  not,  as 
a  rule,  exist  in  the  uncombined  state.  As  soon  as  atoms 
are  freed  from  combination,  they  at  once  unite  with  some 
other  atom  or  atoms.  The  smallest  particle  of  matter 
which  can  exist  independently  is  not,  therefore,  an  atom, 
but  a  group  or  combination  of  atoms.  These  groups  of 
atoms  are  called  molecules.  If  the  atoms  in  a  molecule 
are  atoms  of  the  same  element,  then  the  molecule  is  a 
molecule  of  an  element;  but  if  the  atoms  of  different 
elements  are  combined,  then  the  molecule  is  the  molecule 
of  a  compound.  All  matter,  as  a  rule,  consists  of  mole- 


Chemical  Symbols.  8 

cules,  and  the  molecules  are  made  up  of  atoms.  A  mole- 
cule of  a  few  elements  contains  only  one  atom.  Chemists 
define  a  molecule  as  the  smallest  part  of  a  compound  or 
of  an  element  which  can  exist  in  the  free  state  and  mani- 
fest the  properties  of  the  compound.  Thus,  the  smallest 
particle  of  water  is  a  molecule  of  water,  but  a  molecule  of 
water  contains  smaller  particles  still,  viz.,  atoms  of  hydro- 
gen and  oxygen.  We  may  define  an  atom  as  the  indivis- 
ible constituent  of  a  molecule.  It  is  also  the  smallest 
particle  of  an  element  which  takes  part  in  chemical 
changes. 

Our  views  regarding  molecules  are  based  on  extensive  study  of  the 
physical  properties  of  gases.  The  molecule  is  often  spoken  of  as  the 
physical  unit,  because  in  physical  changes  molecules  are  not  decomposed. 
Whereas  the  atom  is  the  chemical  unit,  because  it  enters  into  all  chemi- 
cal action.  The  molecule  is  chemically  divisible,  but  the  atom  is 
chemically  indivisible. 

Chemical  Symbols,  which  were  mentioned  in  Chapter  I, 
are  designed  to  represent  single  atoms.  Thus,  H  repre- 
sents one  atom  of  hydrogen,  O  one  atom  of  oxygen,  N  one 
atom  of  nitrogen.  If  more  than  one  atom  is  to  be  desig- 
nated, the  proper  numeral  is  placed  before  the  symbol, 

2  H  means  2  atoms  of  hydrogen. 

3  O  means  3  atoms  of  oxygen. 

4  P  means  4  atoms  of  phosphorus. 

But  if  the  atoms  are  in  chemical  combination,  either  with 
themselves  or  with  other  atoms,  then  a  small  numeral  is 
placed  after  and  a  little  below  the  symbol,  thus :  — 

H2  means  2  atoms  of  hydrogen  in  combination, 
N3  means  3  atoms  of  nitrogen  in  combination, 
P4  means  4  atoms  of  phosphorus  in  combination. 


8  2  Descriptive  Chemistry. 

Chemical  Formulas.  —  A  formula  is  a  group  of  symbols 
which  is  designed  to  express  the  composition  of  a  com- 
pound. In  writing  a  formula  the  symbols  of  the  different 
atoms  making  up  the  compound  are  placed  side  by  side. 
Thus,  H2O  is  the  formula  of  water,  because  this  group  of 
symbols  is  the  simplest  expression  of  the  facts  which  are 
known  about  this  compound.  Similarly,  KC1O3  is  the 
formula  of  potassium  chlorate.  These  symbols  might  be 
written  in  a  different  order,  but  usage  has  determined  the 
order  in  this,  as  in  most  cases.  A  formula  represents  one 
molecule.  Hence,  KC1O3  represents  one  molecule  of 
potassium  chlorate,  and  means  that  the  molecule  of  this 
compound  contains  one  atom  each  of  potassium  and  chlo- 
rine and  three  atoms  of  oxygen.  If  we  wish  to  designate 
several  molecules,  the  proper  numeral  is  placed  before  the 
formula,  thus :  — 

2  KC1O3  means  2  molecules  of  potassium  chlorate. 

3  H2O      means  3  molecules  of  water. 

4  H2SO4  means  4  molecules  of  sulphuric  acid. 

In  certain  compounds  some  of  the  atoms  act  like  a  single 
atom  in  chemical  changes.  This  fact  is  often  expressed  by 
inclosing  the  group  of  atoms  in  a  parenthesis,  or  by  sepa- 
rating it  from  the  rest  of  the  formula  by  a  period.  Thus, 
the  formula  of  ammonium  nitrate  is  (NH4)NO3.  Simi- 
larly, the  formula  of  alcohol  is  often  written  C2H5 .  OH, 
because  the  groups  C2H5  and  OH  act  as  units.  The  use 
of  the  period  is  confined  mainly  to  organic  and  mineralogi- 
cal  chemistry.  It  is  sometimes  omitted,  especially  if  the 
composition  of  the  compound  is  well  understood.  If  a 
group  of  atoms  is  to  be  multiplied,  it  is  placed  within  a 
parenthesis.  Thus,  the  formula  of  lead  nitrate  is  Pb(NO3)2. 
This  means  that  the  group  NO3  is  to  be  multiplied  by  2. 


Chemical   Equations.  83 

The  formula  2  Pb(NO3)2  means  that  the  whole  formula  is 
to  be  multiplied  by  2. 

Symbols  and  formulas  are  sometimes  used  to  represent  an  indefinite 
amount  of  an  element  or  compound.  Thus,  O  may  mean  oxygen  and 
H.jSO4  sulphuric  acid,  regardless  of  the  amount.  This  use  of  symbols 
and  formulas  saves  time,  but  it  is  not  scientific.  They  are  often  thus 
used  to  label  bottles  in  a  laboratory.  Such  a  departure  from  accuracy 
should  not  be  allowed  to  obscure  their  real  meaning. 

The  complete  significance  ot  symbols  and  formulas  can  be  grasped 
only  by  their  intelligent  use.  They  should  not  be  committed  to  mem- 
ory slavishly.  It  is  desirable,  however,  to  learn  the  common  ones 
while  the  substances  they  represent  are  being  studied,  and  consider 
their  relations  more  fully  when  the  needed  facts  have  accumulated. 
(See  Chapters  IX  and  XIII.) 

A  Chemical  Reaction  is  a  special  or  limited  chemical 
change.  When  potassium  chlorate  is  heated,  the  chemical 
change  results  finally  in  the  liberation  of  all  the  oxygen 
and  the  formation  of  potassium  chloride.  Such  a  change 
is  called  the  reaction  for  preparing  oxygen  from  potassium 
chlorate,  or  the  reaction  for  the  decomposition  of  potas- 
sium chlorate.  Obviously,  the  study  of  chemistry  is 
largely  a  study  of  reactions. 

Chemical  Equations.  —  In  expressing  various  facts 
about  chemical  reactions,  it  is  customary  to  use  an  equa- 
tion consisting  of  the  proper  symbols  or  formulas.  Sub- 
stances entering  into  the  initial  stage  of  a  reaction  are 
called  factors,  and  those  present  in  the  final  stage  are 
called  products.  The  symbols  and  formulas  of  the  factors 
connected  by  the  sign  plus  (  -f- )  are  placed  at  the  left  of 
the  sign  of  equality,  and  those  of  the  products  at  the  right. 
Equations  are  usually  read  from  left  to  right.  Occasion- 
ally the  words  reaction  and  equation  are  used  as  synonyms, 
but  such  a  use  is  inaccurate  and  confusing. 


84  Descriptive  Chemistry. 

When  magnesium  burns  in  the  air  or  in  oxygen,  mag- 
nesium oxide  is  formed.  The  simplest  equation  for  this 
reaction  is  — 

Mg       +   O     =          MgO 
Magnesium  '  Oxygen     Magnesium  Oxide 

This  equation  is  read  :  Magnesium  and  oxygen  form  mag- 
nesium oxide.  It  means,  also,  that  when  magnesium  and 
oxygen' react,  one  atom  of  magnesium  unites  with  one 
atom  of  oxygen  and  forms  one  molecule  of  magnesium 
oxide.  The  simplest  equation  for  the  preparation  of  hy- 
drogen by  the  reaction  of  zinc  and  sulphuric  acid  is  — 

Zn+      H2SO4     =      H2      +     ZnSO4 
Zinc     Sulphuric  Acid     Hydrogen     Zinc  Sulphate 

This  equation  is  read:  Zinc  and  sulphuric  acid  form  (or 
produce)  hydrogen  and  zinc  sulphate.  It  means,  further, 
that  one  atom  of  zinc  and  one  molecule  of  sulphuric  acid 
form  one  molecule  (or  two  atoms)  of  hydrogen  and  one 
molecule  of  zinc  sulphate.  By  similar  equations  we  may 
express  certain  facts  about  all  reactions  which  are  under- 
stood. The  above  equations  might  be  called  ordinary 
chemical  equations,  or  atomic  equations.  Other  forms 
are  used,  and  they  will  be  discussed  in  Chapters  IX,  X, 
and  XIII. 

The  following  facts  about  ordinary  chemical  equations  should  be 
noted :  — 

(1)  The  sign  plus  does  not  necessarily  mean  addition  chemically. 
It  does  in  the  equation  Mg  -f  O  =  MgO,  but  not  in  the  equation  HgO 
=  Hg-fO.     In  the  latter  the  products  are  merely  mixed.     The  sign 
plus  may  be  expressed  by  the  words  and*  acted  upon,  added  to,  mixed 
with.     The  sign  equality  is  often  read  equal,  give,  form,  or  produce. 

(2)  Equations  do  not  always  include  all  the  participating  substances. 
In  Mg  +  O  =  MgO  no  nitrogen  (N)  appears  because  nitrogen  takes  no 


Exercises.  85 

chemical  part  in  the  change,  despite  the  fact  that  the  air  is  largely 
nitrogen.  Similarly,  in  Zn  +  H2SO4  =  H2  +  ZnSO4,  no  water  (H2O) 
appears,  because  the  water  (in  the  dilute  sulphuric  acid)  simply  serves 
to  dissolve  the  zinc  sulphate  from  the  surface  of  the  zinc.  A  special 
form  of  equation,  called  the  ionic  equation,  is  used  to  express  chemical 
changes  which  occur  in  solution  (see  Chapter  X). 

(3)  Equations  tell  nothing  about  the  heat  changes  (see  Chapter  X). 

(4)  Most  equations  represent  only  the  beginning  and  end  of  reac- 
tions.    Thus,  in   KC1O3  =  O3  +  KC1  several   changes   do   not  appear, 
because  the  purpose  of  this  equation  is  to  express  the  complete  decom- 
position of  potassium  chlorate  —  nothing  else. 


EXERCISES. 

1.  Define  law,  theory,  and  hypothesis  as  used  in  science. 

2.  State  the  law  of  definite  proportions.    Illustrate  it.    Give  a  brief 
account  of  its  discovery. 

3.  State  the  law  of  multiple  proportions.     Illustrate  it.     Who  dis- 
covered it?    When? 

4.  State  the  atomic  theory.     What  are  atoms  according  to  this 
theory?     How  are  atoms  related  to  chemical  action?     How  are  atoms 
related  to  molecules?     What  is  a  molecule? 

5.  What  is  the  symbol  of  an  element?     How  are  they  formed? 
Interpret  the  symbols :  H,  2O,  N3,   2  P,   30,   K2,  S2,  2  Cl. 

6.  What  is  the  formula  of  a  compound?    What  does  a  formula 
represent?      Interpret    the   formulas:  H2O,    2  H2O,    KC1O3,   4  H2SO4, 
(NH4)NO3,   C2H5.OH,   Pb(N(X)2,   Ca(OH)2.  " 

7.  Give  the  symbols  of  the  following  elements  :  oxygen,  hydrogen, 
nitrogen,  zinc,  copper,   magnesium,    platinum,  iron,  sodium,  sulphur, 
carbon,  mercury. 

8.  What  elements  correspond  to  the  following  symbols :    Na,  Cu, 
K,  Zn,  S,  P,  Pt,  Pb,  H,  Hg,  Fe,  Mg? 

9.  Give  the  formulas  of  the  following  compounds :   water,  potas- 
sium chlorate,  sulphuric  acid,  magnesium  oxide. 

10.  Define  and  illustrate  the  term  chemical  reaction. 

11.  What  is  a  chemical  equation ?     For  what  is  it  used?     What  are 
factors  and  products  in  an   equation?     How  are  equations   written? 
Illustrate  your  answer.     How  are  they  read  ? 


86  Descriptive  Chemistry. 

12.  Interpret  the  equation  :    Mg  +  O  =  MgO. 

13.  What  does  the  plus   (  +  )   sign  mean  in  the  above  equation? 
What  other  meanings  has  this  sign? 

14.  State  several  facts  about  equations. 


PROBLEMS. 

1.  How  many  centigrams  in  1745  kg.?     In  250  gm.?    In  1425  dg.  ? 

2.  How  many  cubic  centimeters  in  50  1.  ?     In  I  cu.  dm.  ? 

3.  What  is  the  weight  of  (a)  loocc.  of  hydrogen,  and  (6)  25  1.  of 
oxygen,  under  standard  conditions  ? 

4.  What  weight  of  (a)  hydrogen  and  (<£)  oxygen  can  be  obtained 
from  1 80  gm.  of  water  ? 

5.  What  (#)  weight  and  ($)  volume  of  oxygen  are  necessary  to  unite 
with  200  kg.  of  hydrogen  ? 

6.  What  weight  of  hydrogen  is  necessary  to  unite  with  the  oxygen 
in  100  gm.  of  air  to  form  water  ?  (Assume  that  air  is  one  fifth  oxygen.) 


CHAPTER  VIII. 
ACIDS,   BASES,   AND  SALTS. 

Introduction.  —  Many  chemical  compounds  fall  naturally 
into  one  of  three  groups,  long  known  as  acids,  bases,  and 
salts.  Not  all  compounds,  of  course,  are  included  in  this 
classification.  Each  group  has  its  characteristic  properties, 
'though  the  groups  are  closely  related  and  sometimes  over- 
lap. Many  familiar  substances  belong  to  these  groups. 
A  knowledge  of  the  properties  of  acids,  bases,  and  salts,, 
of  their  special  behavior,  and  of  their  intimate  relations  is 
essential  in  the  study  of  chemistry. 

General  Properties  of  Acids,  Bases,  and  Salts.  —  Acids 
have  a  sour  taste.  The  early  chemists  detected  this 
property,  and  the  word  acid  (from  the  Latin  acidtts,  sour) 
emphasizes  the  fact.  Acids  change  the  color  of  many 
vegetable  substances.  Thus,  blue  litmus  is  turned  red  by 
acids.  Acids  also  have  the  power  to  decompose  most 
carbonates,  like  limestone,  thereby  liberating  carbon  diox- 
ide gas  which  escapes  with  effervescence.  Most  bases 
have  a  slimy,  soapy  feeling,  and  a  bitter  taste.  They  turn 
red  litmus  blue.  Caustic  soda  and  ammonium  hydroxide 
are  bases.  Many  salts  have  the  well-known  salty  taste. 
Sodium  chloride,  the  familiar  table  salt,  is  an  example. 
Usually,  they  have  no  action  on  litmus. 

All  acids  contain  hydrogen,  which  is  usually  liberated 
when  metals  and  acids  interact.  Most  acids  contain  oxy- 
gen. For  many  years  it  was  thought  that  oxygen  was  an 

87 


88  Descriptive  Chemistry. 

essential  component  of  all  acids,  and  its  name,  oxygen 
(derived  from  Greek  words  meaning  "  acid  producer  ")  was 
given  by  Lavoisier  because  of  this  belief  (see  Discovery  of 
Oxygen). 

We  now  know  that  hydrogen,  not  oxygen,  is  the 
essential  component  of  all  acids.  Another  necessary 
component  of  acids  is  some  element  like  nitrogen,  sulphur, 
chlorine,  or  phosphorus,  which  belongs  to  a  class  of 
elements  called  non-metals.  For  this  reason  it  is  some- 
times convenient  to  think  of  non-metals  as  the  elements 
which  form  acids.  Thus  sulphuric  acid  contains  sulphur, 
besides  hydrogen  and  oxygen ;  while  hydrochloric  acid 
contains  chlorine,  besides  hydrogen. 

Bases  contain  oxygen  and  usually  hydrogen,  but  their 
distinctive  component  is  a  metal,  e.g.  sodium,  potassium, 
calcium.  Hence  a  metal  may  be  properly  regarded  not 
merely  as  an  element  possessing  in  a  varying  degree  the 
physical  properties  of  hardness,  luster,  power  to  conduct 
heat  and  electricity,  but  also  the  chemical  property  of 
forming  bases. 

Salts  contain  a  metal  and  a  non-metal,  and  most  of  them 
contain  oxygen.  Thus,  potassium  nitrate  contains  the 
metal  potassium  and  the  non-metal  nitrogen,  besides 
oxygen ;  while  potassium  chloride  contains  potassium 
and  the  non-metal  chlorine,  but  no  oxygen. 

The  nature  o*f  acids,  bases,  and  salts  is  clearly  shown  by 
their  chemical  relations  to  each  other.  When  acids  and 
bases  interact,  salts  are  formed.  That  is,  the  acid  and 
base  destroy  more  or  less  completely  the  marked  prop- 
erties of  each  other  and  produce  a  compound  which  has 
few,  and  often  none,  of  the  properties  of  the  original  acid 
or  base.  The  acid  and  base  neutralize  each  other.  An 
example  will  make  this  point  clear.  When  hydrochloric 


Acids,   Bases,  and  Salts.  89 

acid  and  sodium  hydroxide  interact,  sodium  chloride  and 
water  are  formed.  The  chemical  change  may  be  written 

thus- 

HC1          +  NaOH.  NaCl      +      H2O 

Hydrochloric  Acid       Sodium  Hydroxide       Sodium  Chloride       Water 

This  equation  represents  the  facts  which  have  been 
repeatedly  verified  by  experiment.  This  series  of  chemi- 
cal changes  is  called  neutralization,  and  later  it  will  be 
more  fully  discussed.  Taking  this  equation  as  a  type  of 
the  chemical  changes  which  occur  in  neutralization,  it  is 
clear  that  in  such  changes,  generally  speaking  (i)  the  metal 
of  the  base  takes  the  place  of  the  hydrogen  of  the  acid, 
thereby  forming  a  salt,  while  (2)  the  hydrogen  of  the  acid 
combines  with  the  hydrogen  and  oxygen  of  the  base  to 
form  water.  In  neutralization  the  hydrogen  and  oxygen 
of  the  base  act  as  a  unit.  This  group  of  atoms  (OH)  is 
called  hydroxyl.  Compounds  containing  this  group  are 
called  hydroxides.  Hydroxyl  does  not  exist  free  and 
uncombined  like  elements  and  compounds,  but  it  acts  like 
a  single  atom  in  many  changes.  It  is  called  a  radical. 
To  emphasize  the  fact  that  it  is  a  unit,  the  hydroxyl  group 
is  sometimes  put  in  a  parenthesis,  e.g.  Ca(OH)2. 

Hydroxides  are  often  said  to  be  founded  on  the  water  type.  Thus 
we  have  — 

Water HOH 

Sodium  hydroxide         .         .         .         .   '     NaOH 
Potassium  hydroxide    ....         KOH 
Calcium  hydroxide        ....         Ca(OH)2 

Hence  we  may  regard  sodium  hydroxide  and  potassium  hydroxide 
as  water  in  which  the  hydrogen  atom  has  been  replaced  by  a  metallic 
atom. 

The  words  hydroxide,  hydrate,  and  hydroxyl  are  all  derived  from 
hudor,  the  Greek  word  for  water. 


90  Descriptive  Chemistry. 

The  most  characteristic  property  of  acids  and  bases  is, 
then,  this  power  to  neutralize  each  other  and  thereby  form 
salts  and  water. 

Acids.  —  The  common  acids  are  sulphuric  acid,  hydro- 
chloric acid,  nitric  acid,  and  acetic  acid.  Many  acids  are 
liquid,  as  sulphuric  and  nitric ;  a  few  are  gases,  as  hydro- 
chloric ;  others  are  solid,  as  tartaric,  citric,  oxalic.  Most 
are  soluble  in  water,  and  such  solutions  are  familiarly 
called  acids.  These  solutions  may  be  dilute  or  concen- 
trated, and  the  general  properties  vary  somewhat  with  the 
strength.  Concentrated  acids  are  usually  corrosive  and 
should  be  handled  with  precaution,  even  when  one  is 
thoroughly  familiar  with  their  properties.  Substances 
which  turn  blue  litmus  to  red  are  said  to  contain  an  acid, 
to  be  acid,  or  to  have  an  acid  reaction.  The  exact  nature, 
however,  of  such  a  substance  must  be  determined  by 
additional  tests. 

Many  familiar  substances  are  acids  or  contain  them. 
Vinegar,  pickles,  and  similar  relishes  contain  dilute  acetic 
acid.  Lemon  juice  is  mainly  citric  acid.  Sour  milk  con- 
tains lactic  acid.  Unripe  fruits,  sour  bread,  and  sour 
wines  contain  acids.  "  Soda  water "  is  a  solution  of 
carbonic  acid  (or  more  accurately  carbon  dioxide),  and 
"  acid  phosphate"  is  a  solution  of  a  sour  calcium  phosphate. 

No  brief,  satisfactory  definition  of  an  acid  can  be  given, 
for  chemists  do  not  agree  on  this  point.  We  might  say, 
however,  that  an  acid  is  a  compound  containing  hydrogen 
which  can  be  replaced  by  a  metal;  but  this  definition 
includes  water,  since  its  hydrogen  is  readily  replaced  by 
sodium.  Not  only  must  the  hydrogen  of  an  acid  be 
replaced  by  a  metal,  but  one  product  of  the  reaction  must 
be  a  salt.  The  replacing  metal  may,  of  course,  come  from 
a  compound,  e.g.  an  oxide,  hydroxide,  or  carbonate. 


Acids,  Bases,  and  Salts.  91 

Nomenclature  of  Acids.  —  Oxygen  is  a  component  of 
most  acids,  and  the  names  of  these  acids  correspond  to 
the  proportion  of  oxygen  which  they  contain.  The  best 
known  acid  of  an  element  usually  has  the  suffix  -ic,  e.g. 
sulphuric,  nitric,  phosphoric.  If  an  element  forms  another 
acid,  containing  less  oxygen,  this  acid  has  the  suffix  -ous, 
e.g.  sulphurous,  chlorous,  phosphorous.  Some  elements 
form  an  acid  containing  less  oxygen  than  the  -ous  acid ; 
these  acids  retain  the  suffix  -ous,  and  have,  also,  the  prefix 
hypo-,  e.g.  hyposulphurous,  hypophosphorous,  hypochlo- 
rodl.  Hypo-  means  under  or  lesser.  If  an  element  forms 
an  acid  containing  more  oxygen  than  the  -ic  acid,  such  an 
acid  retains  the  suffix  -ic,  and  has,  also,  the  prefix  per-,  e.g. 
persulphuric,  perchloric.  The  prefix  per-  means  beyond 
or  over.  The  few  acids  which  contain  no  oxygen  have 
the  prefix  hydro-  and  the  suffix  -ic,  e.g.  hydrochloric, 
hydrobron\i£,  hydrofluoric.  It  should  be  noticed  that 
these  suffixe^  are  not  always  added  to  the  name  of  the 
element,  but  often  to  some  modification  of  it. 

The  nomenclature  of  acids  is  well  illustrated  by  the  series  of  chlorine 
acids :  — 

*  ACIDS  OF  THE  ELEMENT  CHLORINE. 


NAME. 

FORMULA. 

Hydrochloric 
Hypochlorous 
Chlorous 

HC1 
HC1O 
HC1O2 

Chloric 
Perchloric 

HC103 
HC1O4 

Not  all  elements  form  a  complete  series  of  acids,  but  the 
nomenclature   usually  agrees  with   the   above   principles. 


92  Descriptive  Chemistry. 

Some  acids  have  commercial  names.  Thus,  sulphuric  acid 
is  often  called  oil  of  vitriol,  and  hydrochloric  acid  is  known 
as  muriatic  acid.  Acids  in  which  carbon  is  the  essential 
component  end  hi  -ic,  but  they  are  often  arbitrarily  named 
(see  Organic  Acids). 

An  examination  of  the  formulas  of  acids  shows  that  all  do  not  con- 
tain the  same  number  of  hydrogen  atoms.  Acids  are  sometimes  classi- 
fied by  the  number  of  hydrogen  atoms  which  can  be  replaced  by  a  metal. 
This  varying  power  of  replaceability  is  called  basicity.  A  monobasic 
acid  contains  only  one  atom  of  replaceable  hydrogen  in  a  molecule,  e.g. 
nitric  acid,  HNO3.  A  molecule  of  acetic  acid  (C2H4O2)  contains  four 
atoms  of  hydrogen,  but  for  reasons  which  are  too  complex  to  state  here, 
only  one  of  these  atoms  can  be  replaced  by  a  metal.  Dibasic  and 
tribasic  acids  contain  two  and  three  replaceable  hydrogen  atoms,  e.g. 
sulphuric  acid  (H2SO4)  and  phosphoric  acid  (H3PO4).  Obviously, 
monobasic  acids  form  only  one  class  of  salts,  dibasic  acids  form  two 
classes,  tribasic  acids  form  three,  and  so  on. 

Bases.  —  The  term  base,  in  a  narrow  sense,  means  the 
strong  bases,  which  are  very  soluble  in  water,  and  are  com- 
monly known  as  alkalies,  e.g.  sodium,  potassium,  and 
ammonium  hydroxides.  In  a  broad  sense  it  means  any 
substance  which  will  neutralize  an  acid,  e.jr.  calcium  oxide, 
ammonia  gas,  as  well  as  the  hydroxides  of  metals.  Most 
bases  are  solids ;  but  since  they  are  usually  soluble  in  water, 
these  solutions,  as  in  the  case  of  acids,  are  familiarly 
called  the  base,  or  alkali,  itself.  Concentrated  alkalies, 
like  concentrated  acids,  are  corrosive.  The  common  alka- 
lies —  sodium  and  potassium  hydroxides  —  are  often  called 
caustic  soda  and  caustic  potash  to  emphasize  this  property  ; 
and  calcium  oxide,  or  lime,  is  sometimes  called  caustic 
lime ;  the  corrosive  nature  of  ammonium  hydroxide,  or 
ordinary 'ammonia,  is  also  well  known.  Substances  which 
turn  red  litmus  to  blue  are  said  to  contain  an  alkali  (or 
base),  to  be  alkaline,  or  to  have  an  alkaline  reaction. 


Acids,  Bases,  and  Salts.  93 

The  word  basic  is  often  used  instead  of  alkaline.  Other 
tests  besides  that  with  litmus  must  be  applied,  however,  to 
determine  the  exact  nature  of  a  substance  having  an  alka- 
line reaction.  Alkalies  dissolve  grease  and  fats,  and  are 
often  used  as  cleansing  agents,  ammonium  hydroxide 
being  widely  employed  for  this  purpose.  They  also  inter- 
act with  fats  to  form  soaps,  large  quantities  of  sodium 
hydroxide  being  annually  utilized  in  the  soap  industry  (see 
Soap). 

A  base,  like  an  acid,  is  rather  difficult  to  define.  We 
might  say  that  a  base  is  an  hydroxide  or  oxide  of  a  metal, 
which  will  neutralize  an  acid,  thereby  forming  a  salt. 
The  term  must  include  ammonia,  which  does  not  contain 
a  metal.  But,  as  we  shall  see  later,  a  certain  combination 
of  elements  related  to  ammonia  acts  like  a  metal  (see 
Ammonium). 

Nomenclature  of  Bases.  —  There  is  no  general  rule 
covering  the  nomenclature  of  bases,  as  in  the  case  of 
acids.  Since  most  bases  contain  hydrogen  and  oxygen, 
they  are  often  called  hydroxides.  Hydrate  is  sometimes 
used  as  a  synonym  of  hydroxide.  The  term  alkali  em- 
phasizes general  properties  rather  than  suggests  specific 
composition.  Hydroxides  are  distinguished  from  each 
other  by  placing  the  name  of  the  metal  before  the  word 
hydroxide,  e.g.  sodium  hydroxide,  potassium  hydroxide, 
calcium  hydroxide.  The  common  hydroxides  have  long 
been  known  by  several  names.  Thus,  calcium  hydroxide 
is  often  called  limewater.  Ammonium  hydroxide  is  some- 
times called  ammonia  water  or  simply  (but  inaccurately) 
ammonia,  and  it  was  formerly  called  volatile  alkali.  Be- 
sides the  common  names  of  the  hydroxides  of  sodium  and 
potassium  already  given,  they  are  sometimes  called  fixed 
alkalies. 


94  Descriptive  Chemistry. 

Not  all  bases  contain  the  same  number  of  hydroxyl  groups.  Hence 
bases,  like  acids,  may  form  one  or  more  salts.  This  power  is  called 
acidity.  Bases  are  called  monacid,  diacid,  triacid  bases,  etc.,  accord- 
ing to  the  number  of  replaceable  hydroxyl  groups  present  in  a  molecule. 
Thus,  calcium  hydroxide  (Ca(OH)2)  is  a  diacid  base,  and  aluminium 
hydroxide  (A1(OH)3)  is  a  triacid  base. 

Salts.  —  Sodium  chloride,  or  ordinary  table  salt,  is  the 
most  familiar  salt.  It  has  been  known  for  ages.  Doubt- 
less this  class  of  chemical  compounds  received  its  name 
because  of  the  general  resemblance  most  of  them  bear  to 
common  salt.  Most  salts  are  solid  and  are  soluble  in 
water.  Many  of  them  have  no  action  on  litmus,  and  are, 
therefore,  said  to  be  neutral  or  to  have  a  neutral  reaction. 
This  indifference  to  litmus  is  not  a  decisive  test  for  a 
salt,  since  many  other  substances,  water  for  example, 
have  no  action  on  litmus.  Nevertheless  the  term  neutral 
is  applied  to  substances  which  do  not  change  the  color 
of  litmus. 

Some  substances  which  are  salts,  as  far  as  their  structure  and  method 
of  formation  are  concerned,  do  not  have  a  neutral  reaction.  Thus, 
sodium  carbonate,  which  is  the  sodium  salt  of  carbonic  acid,  has  a 
marked  alkaline  reaction,  being  in  fact  known  in  commerce  simply  as 
"  alkali." 

A  salt  may  be  defined  as  the  main  product  of  the  inter- 
action of  an  acid  and  a  base.  It  may,  however,  be  a  sub- 
stance which  has  the  properties  of  a  salt,  regardless  of  the 
method  of  formation. 

Salts  are  formed  in  various  ways.  The  interaction  of  an  acid  and  a 
base  has  been  mentioned.  The  interaction  of  acids  with  oxides  of  cer- 
tain metals  or  with  metals  themselves  produces  salts.  Sodium  oxide 
and  sulphuric  acid  interact  and  form  the  salt  sodium  sulphate,  thus :  — 

Na,O          +  H2SO4  Na2S04        +        H2O 

Sodium  Oxide         Sulphuric  Acid         Sodium  Sulphate        Water 


Acids,   Bases,  and  Salts.  95 

While  zinc  and  sulphuric  acid,  as  already  stated,  form   the  salt  zinc 
sulphate  as  well  as  hydrogen,  thus  :  — 

Zn        +        H2S04        =  ZnS04       +  H2 

Zinc        Sulphuric  Acid        Zinc  Sulphate        Hydrogen 

Carbonates  interact  with  acids  and  form  other  salts.    Calcium  carbonate 
and  hydrochloric  acid  form  the  salt  calcium  chloride,  thus  :  — 

CaC03       +       2HC1        =        CaCl2      +      CO2     +      H2O 
Calcium         Hydrochloric          Calcium          Carbon         Water 
Carbonate  Acid  Chloride         Dioxide 

Nomenclature  of  Salts.  —  The  name  of  salts  containing 
oxygen  are  derived  from  the  name  of  the  corresponding 
acid.  The  characteristic  suffix  of  the  acid  is  changed  to 
indicate  this  relation.  Thus,  the  suffix  -ic  becomes  -ate, 
and  the  suffix  -ous,  becomes  -ite.  Hence  :  — 

Sulphuric  acid  forms  sulphates. 
Sulphurous  acid  forms  sulphites. 
Nitric  acid  forms  nitrates. 
Nitrous  acid  forms  nitrites. 
Chloric  acid  forms  chlorates. 
Hypochlorous  acid  forms  hypochlorites. 
Permanganic  acid  forms  permanganates.       $ 

The  name  of  the  replacing  metal  is  retained,  e.g.  potas- 
sium chlorate,  sodium  sulphate,  calcium  hypochlorite,  po- 
tassium permanganate.  Notice  that  the  prefixes  hypo-  and 
per-  are  not  changed. 

The  names  of  salts  containing  only  two  elements,  fol- 
lowing the  general  rule  for  binary  compounds,  end  in  -ide. 
This  suffix  is  added  to  a  modification  of  the  name  of  the 
non-metal,  giving  the  names  chloride,  bromide,  sulphide, 
fluoride,  etc.  The  prefix  hydro-  which  is  contained  in  the 


96  Descriptive  Chemistry. 

name  of  the  acid  is  omitted.  Thus,  the  name  of  the 
sodium  salt  of  hydrochloric  acid  is  sodium  chloride ;  simi- 
larly, there  are  the  names  potassium  chloride,  calcium 
fluoride,  and  sodium  iodide.  Sometimes,  the  salts  of  these 
hydrogen  acids  are  called  halides  to  emphasize  their  rela- 
tion to  common  salt,  which  in  Greek  is  called  halos. 

Salts  in  which  all  the  hydrogen  atoms  of  the  corresponding  acid 
have  been  replaced  by  a  metal  are  called  normal  salts,  e.g.  sodium 
sulphate,  Na.,SO4.  If  some  of  the  hydrogen  atoms  are  not  replaced  by 
a  metal,  an  acid  salt  is  formed.  Thus,  acid  sodium  sulphate  may  be 
regarded  as  derived  from  sulphuric  acid,  which  is  dibasic,  by  replacing 
one  of  the  atoms  of  hydrogen  by  sodium,  though  of  course  the  salt  is 
not  prepared  in  this  way.  Expressed  as  formulas  these  relations  may 
be  written  thus  :  — 

Acid  Acid  Salt  Normal  Salt 

H,SO4  HNaSO4  Na2SO4 

Only  those  acids  which  contain  two  or  more  replaceable  hydrogen 
atoms  form  acid  salts.  On  the  other  hand,  if  not  all  the  hydroxyl 
groups  of  a  base  are  replaced  when  the  base  reacts  with  an  acid,  then  a 
basic  salt  results.  Thus,  basic  nitrate  of  bismuth  may  be  regarded  as 
the  salt  derived  from  bismuth  hydroxide  (Bi(OH)3)  by  replacing  one 
hydroxyl  group  of  the  base  by  the  group  NO3  of  nitric  acid.  The 
formula  of  this  basic  nitrate  of  bismuth  is  Bi(OH)2NO3. 

The  following  equation  illustrates  the  changes  :  — 

Bi(OH)3         +  HN03        =        Bi(OH)2NOo       +        H2O 

Bismuth  Hydroxide        Nitric  Acid        Basic  Bismuth  Nitrate       Water 

Only  those  bases  having  two  or  more  hydroxyl  groups  can  form  basic 
salts.  Some  basic  salts  are  very  complex. 

Relation  of  Oxides  to  Acids  and  Bases.  —  Most  non- 
metallic  elements  form  oxides  which  unite  with  water  and 
produce  an  acid.  The  oxides  of  many  metallic  elements, 


Acids,  Bases,  and  Salts.  97 

on  the  other  hand,  unite  with  water  and  produce  hydrox- 
ides.    The  two  oxides  of  the  non-metal  sulphur  act  thus  — 

502  +       H2O      =      H2SO3 
Sulphur  dioxide        Water          Sulphurous  Acid 

503  +        H2O      =     H2SO4 
Sulphur  Trioxide        Water          Sulphuric  Acid 

The  oxide  of  the  metal  calcium  acts  thus  — 

CaO       -f       H2O      =      Ca(OH)2 
Calcium  Oxide        Water          Calcium  Hydroxide 

Oxides  of  non-metals  which  unite  with  water  and  thereby 
produce  acids  are  called  anhydrides,  i.e.  literally,  sub- 
stances without  water.  Examples  are  carbonic  anhydride 
(CO2),  sulphuric  anhydride  (SO3),  phosphoric  anhydride 
(P2O5).  Oxides  of  metals  which  produce  hydroxides  are 
called  basic  oxides.  A  few  oxides  behave  exceptionally. 
It  is  convenient  to  regard  an  anhydride  as  the  root  or 
basis  of  its  corresponding  acid,  and  a  basic  oxide  as  the 
root  of  its  hydroxide. 

The  fact  that  many  non-metallic  oxides  redden  moist  blue  litmus  led 
Lavoisier  into  the  erroneous  belief  that  oxygen  is  an  essential  compo- 
nent of  acids.  And  some  authorities  even  now  (incorrectly)  speak  of 
these  oxides  as  acids  ;  thus,  carbon  dioxide  (CO2)  is  occasionally  called 
carbonic  acid.  The  compounds  which  Lavoisier  galled  acids  were  anhy- 
drides. And  it  was  not  until  about  181 1  that  Davy  showed  (i)  that  some 
acids  do  not  contain  oxygen  (e.g.  hydrochloric  acid,  HC1 ),  and  (2)  that 
the  so-called  acids  of  Lavoisier  are  not  real  acids  until  they  have  obtained 
hydrogen  from  the  water  with  which  they  combine. 

Neutralization  has  been  defined  as  the  series  of  changes 
whereby  acids  and  bases  mutually  destroy  each  other's 
characteristic  properties  and  produce  a  salt  and  water. 


98 


Descriptive  Chemistry. 


But  neutralization  has  a  deeper  meaning  and  broader  ap- 
plication than  the  mere  destruction 
of  properties. 

If  measured  volumes  of  different  acids 
are  exactly  neutralized  by  different  alkalies, 
remarkable  relations  are  revealed.  This  may 
be  done  by  dropping  one  into  the  other  from 
a  graduated  tube,  called  a  burette  (Fig.  15). 
The  exact  point  of  neutralization  is  shown 
by  an  indicator;  this  is  a  solution  of  litmus 
or  some  other  substance,  which  tells  by  the 
color  whether  the  solution  is  acid  or  alkaline. 
Experiment  shows  that  (i)  a  definite  quan- 
tity of  an  acid  neutralizes  a  definite  quantity 
of  an  alkali,  (2)  the  same  acid  is  neutralized 
by  different  quantities  of  different  alkalies, 
and  (3)  the  ratio  of  the  quantities  of  the 
FIG.  15.  — Burettes.  different  alkalies  is  the  same  for  all  acids.1 


EXERCISES. 

1.  Define  and  illustrate  (a)  an  acid,  (6)  a  base,  (c)  a  salt,  (d}  an  al- 
kali, (e)  hydroxyl,  (/)  an  hydroxide. 

2.  Name  three  common  acids  and  bases.     State  the  general  proper- 
ties of  each  class. 

3.  Define  and  illustrate  (a)  neutralization,  (b}  acidity  of  bases,  (c} 
basicity  of  acids,  (</)  normal,  acid,  and  basic  salts,  (V)  caustic  alkali, 
(/)  radical. 

4.  What  is  the  literal  meaning  of  (a)  acid  (  adj.),  (£)  caustic,  (c)  per-, 
(d)  hypo-,  (i)  anhydride? 

5.  Name  the  sodium  salt  of  hydrochloric  acid.     Name  the  corre- 
sponding salt  of  potassium,  lead,  calcium,  barium,  zinc,  silver. 

6.  Name  the  same  salts  of  nitric  acid.     Of  nitrous  acid. 

7.  Name  the  same  salts  of  sulphuric  acid.     Of  hypochlorous  acid. 
Of  perchloric  acid. 

1  A  more  extended  treatment  of  this  subject  may  be  found  in  the  author's 
"Experimental  Chemistry,"  pp.  124  ff. 


Acids,   Bases,  and  Salts.  99 

8.  Name  the  hydroxides  corresponding  to  sodium,  potassium,  calcium, 
barium,  zinc,  lead,  copper. 

9.  Name  the  potassium  salt  of  manganic  acid,  calcium  salt  of  hydro- 
fluoric acid,  sodium  salt  of  carbonic  acid,  potassium  salt  of  tartaric  acid, 
lead  salt  of  chromic  acid,  potassium  salt  of  hydrobromic  acid,  potassium 
salt  of  permanganic  acid. 

PROBLEMS. 

Review  any  of  the  preceding  problems,  especially  those  in  Appendix, 
§1. 


CHAPTER   IX. 

EQUIVALENTS  —  ATOMIC  AND  MOLECULAR  WEIGHTS  — 
CHEMICAL  CALCULATIONS— QUANTITATIVE  SIGNIFI- 
CANCE OF  EQUATIONS. 

Equivalents. — The  equivalent  or  equivalent  weight  of 
an  element  is  that  weight  which  is  chemically  equivalent 
to  one  part  by  weight  of  hydrogen.  More  specifically,  it 
is  the  number  of  grams  of  an  element  which  liberates, 
replaces,  or  combines  with  I  gm.  of  hydrogen.  Ex- 
periments show  that  approximately  32.5  gm.  of  zinc  will 
liberate  I  gm.  of  hydrogen  from  an  acid.  Hence  32.5 
is  the  equivalent  of  zinc.  Similarly,  23  gm.  of  sodium 
liberate  I  gm.  of  hydrogen  from  water.  A  summary  of 
numerous  experiments  reveals  the  following  — 

TABLE  OF  EQUIVALENTS. 


ELEMENT. 

EQUIVALENT. 

Hydrogen 

, 

(by  definition) 

Oxygen 
Chlorine 

8 

35-5 

Bromine 

80 

Sulphur 

16 

Zinc 

32-5 

Copper 

3i-7 

Magnesium 

12 

Sodium 

23 

Potassium 

39 

Silver 

108 

Aluminium 

9 

100 


Atomic  '  Weights.  \ "- *." ' ;  \ /,  \\  \ :  /,  i  o  i 

Analysis  of  chemical  compounds  determines  the  propor- 
tion of  their  components  by  weight.  And  in  many  cases 
such  experiments  verify  the  equivalents  found  by  other 
methods.  Thus,  experiment  shows  that  — 

35-5  Par*s  of  chlorine  unite  with  23  of  sodium,  or  39  of  potassium. 

80      parts  of  bromine  unite  with  23  of  sodium,  or  39  of  potassium. 

108      parts  of  silver       replace        23  of  sodium,  or  39  of  potassium. 

The  above  elements  always  unite  in  these  proportions. 
But  some  elements  unite  in  several  proportions.  Thus, 
eight  parts  by  weight  of  oxygen  combine  with  one  part  of 
hydrogen  to  form  water.  But  in  a  large  number  of  com- 
pounds sixteen  parts  of  oxygen  combine  with  various  parts 
of  different  elements.  Similarly,  nitrogen  unites  in  the 
proportion  of  fourteen,  twenty-eight,  and  forty-two  parts 
by  weight  with  different  parts  of  other  elements.  In  a 
word,  there  are  multiples  of  equivalents.  Comparison 
shows  a  striking  coincidence  between  many  equivalent 
weights  and  the  accepted  atomic  weights  of  the  same 
elements.  This  topic  is  discussed  and  applied  in  Chapter 
XIII. 

Atomic  Weights.  —  One  of  the  essential  properties  of 
matter  is  weight.  According  to  the  atomic  theory,  atoms 
have  weight.  But  the  weight  of  an  atom  is  so  small  that 
we  cannot  determine  it.  We  can,  however,  find  the  rela- 
tive weight  of  an  atom ;  that  is,  how  many  times  heavier 
one  atom  is  than  another  atom.  If  we  adopt  one  as  the 
weight  of  an  atom  of  hydrogen,  the  weights  of  atoms  of— 
other  elements  can  be  readily  expressed  in  terms  of  this 
standard.  Thus,  when  we  say  the  atomic  weight  of  sodium 
is  twenty-three,  we  mean  that  an  atom  of  sodium  weighs 
twenty-three  times  as  much  as  an  atom  of  hydrogen.  The 


IO2  Descriptive  Chemistry. 

determination  of  the  exact  atomic  weight  of  an  element  is 
a  difficult  task.  Many  principles  influence  the  final  selec- 
tion of  the  number  adopted  as  the  atomic  weight.  We 
have  already  seen  that  there  is  a  definite  relation  between 
the  equivalent  weight  and  the  atomic  weight  of  an  element. 
But  this  method  cannot  be  used  exclusively  to  determine 
atomic  weights,  because  it  does  not  enable  us  to  tell  the 
number  of  atoms  in  a  molecule.  There  is  also  a  definite 
relation  between  the  molecular  weight  of  a  compound  and 
the  atomic  weights  of  the  elements  in  the  compound. 
These  topics  and  others  related  to  them  will  be  discussed 
in  Chapter  XIII.  For  the  present,  the  approximate  atomic 
weights  found  in  the  Appendix,  §  5,  may  be  used  in  solv- 
ing problems  and  interpreting  equations. 

The  atomic  weights  are  not  necessarily  whole  numbers,  but  they  are 
nearly  so  in  many  cases,  and  for  most  purposes  round  numbers  may  be 
used.  Different  atomic  weights  are  sometimes  given  for  the  same  ele- 
ment. This  is  due  (i)  to  the  disagreement  among  chemists  as  to  the 
accuracy  of  certain  results,  and  (2)  to  the  use  of  several  standards  for 
reckoning  atomic  weights.  For  many  years  hydrogen  was  the  standard. 
But  for  scientific  reasons  oxygen  is  being  adopted  as  the  standard,  and 
1 6  is  accepted  as  its  atomic  weight.  This  change  does  not  alter  the 
facts;  it  merely  changes  the  relative  values  of  the  atomic  weights. 
Thus,  the  atomic  weight  of  hydrogen  becomes  1.008,  if  oxygen  equals 
1 6,  and  others  are  proportionally  changed. 

Tables  of  atomic  weights  have  been  prepared  on  both  standards 
(H  =  i  and  O  =  16).  Both  tables  are  given  in  the  Appendix,  §5. 

Symbols  and  Atomic  Weights.  —  Symbols  not  only  rep- 
resent atoms,  but  they  express  atomic  weights.  Thus,  O 
represents  one  atom  of  oxygen,  but  it  also  means  that  this 
atom  weighs  sixteen  times  more  than  an  atom  of  hydro- 
gen. Similarly,  K  represents  an  atom  of  potassium, 
which  weighs  thirty-nine  times  more  than  an  atom  of 
hydrogen. 


Chemical  Calculations.  103 

Molecular  Weights.  —  Since  atoms  combine  to  form 
molecules,  a  molecular  weight  is  the  sum  of  the  weights 
of  the  atoms  in  a  molecule.  A  molecule  of  nitric  acid 
contains  one  atom  each  of  hydrogen  and  nitrogen,  and 
three  atoms  of  oxygen ;  hence  its  molecular  weight  is 
i  +  14+  16  x  3  =  63.  Given  the  formula,  the  molecular 
weight  is  easily  found  by  adding  the  atomic  weights. 
The  molecular  weight  and  formula  of  a  compound,  there- 
fore, are  rigidly  connected;  and  just  as  a  symbol  stands 
for  an  atomic  weight,  so  a  formula  expresses  a  molecular 
weight.  It  is  customary  to  assume  the  simplest  formula 
(i.e.  the  one  corresponding  to  the  lowest  molecular  weight) 
until  experiments  show  which  is  the  correct  one. 

Many  facts  and  principles  determine  the  final  selection  of  the  molecular 
weight,  and  hence  the  formula,  of  a  compound.  These  will  be  discussed 
in  Chapter  XIII. 

Chemical  Calculations  are  largely  based  on  atomic  and 
molecular  weights. 

Percentage  Composition.  —  Since  the  formula  of  a  com- 
pound expresses  its  composition,  it  is  possible  to  calculate 
from  the  formula  the  composition  in  per  cent.  The  for- 
mula of  sulphuric  acid  is  H2SO4,  and  its  molecular  weight 
is  98,  i.e.  2  +  32+64.  The  calculations  are  most  easily 
made  by  the  following  proportions:  — 

2  :  98  :  :  x  :  100,   x  =    2.04  per  cent  of  hydrogen. 
32  :  98  :  :  x  :  100,   x—  32.65  per  cent  of  sulphur. 
64  :  98  :  :  x  :  100,    x  =  65.31  per  cent  of  oxygen. 
Total  100.00  per  cent. 

By  the  same  method  the  percentage  composition  of  any 
compound  may  be  calculated. 


IO4  Descriptive  Chemistry. 

Simplest  Formula. — The  simplest  formula  of  a  compound 
may  be  found  by  dividing  the  percentage  of  each  element 
in  the  compound  by  its  atomic  weight.  The  percentage 
composition  of  sulphuric  acid  is  H  =  2.04,  8  =  32.65, 
0  =  65.31.  Dividing  each  percentage  by  the  atomic 
weight  of  the  element,  we  have(approximately)2.O4  ^-1=2, 
32.65-^-32=1,  65.31-^16  =  4.  Hence  the  simplest  for- 
mula of  sulphuric  acid  is  H2SO4.  Sometimes  the  prod- 
ucts of  the  percentages  divided  by  the  atomic  weights 
are  not  whole  numbers.  In  that  case  the  simplest  relation 
is  found  by  proportion.  The  following  problem  illustrates 
this  principle :  the  percentage  composition  of  a  compound 
is  C  =  40,  H  =  6.67,  0  =  53.33.  Dividing  as  above,  we 
have  40 H-  12  =  3.33,  H-^  i  =6.67,  53-33^  16=3.33.  But 
3.33,  6.67,  3.33  are  in  the  same  proportion  as  1:2:1. 
Hence  the  simplest  formula  is  CH2O. 

Quantitative  Significance  of  Equations.  —  It  is  possible 
to  express  reactions  in  the  form  of  equations  because  in 
every  chemical  change  no  weight  is  lost  or  gained. 

It  has  already  been  stated  that  the  equation  for  the 
reaction  between  magnesium  and  oxygen  is  — 

Mg       +       O  MgO 

Magnesium          Oxygen       Magnesium  Oxide 

This  equation  is  the  outcome  of  the  following :  it  can  be 
readily  shown  by  experiment  that  when  magnesium  is 
heated  in  air  or  oxygen,  the  magnesium  and  oxygen  com- 
bine in  the  ratio  3  :  2.  Now  results  like  this  are  usually 
expressed  in  terms  of  the  atomic  weights  of  the  reacting 
elements.  But  we  do  not  know  the  number  of  atomic 
weights  of  these  elements  which  must  be  taken  to  produce 
the  ratio  3  :  2.  That  is,  we  do  not  know  whether  the 
ratio  requires  the  atomic  weight  or  some  multiple  of  it. 


Quantitative  Significance  of  Equations.       105 

But,  if  we  let  y  equal  the  unknown  number  of  atomic 
weights  of  magnesium  and  z  the  unknown  number  of 
atomic  weights  of  oxygen,  then  we  can  write  the  prelimi- 
nary equation  thus  — 

y  X  at.  wt.  of  mag.  :  z  X  at.  wt.  oxygen  =  3:2. 

The  atomic  weight  of  magnesium  is  24  and  of  oxygen  is 
1 6.  Therefore  the  problem  reduces  itself  to  finding  the 
values  of  y  and  z  in  the  equation  — 

y  x  24  :  s  x  16  =  3:  2. 

Obviously,  y  =  i  and  z  =  i.  Now  the  symbol  Mg  stands 
for  24  parts  of  magnesium  and  O  for  16  parts  of  oxygen. 
That  is,  Mg  not  only  means  one  atom  of  magnesium,  but 
also  that  this  atom  weighs  24,  if  one  atom  of  oxygen 
weighs  1 6.  Therefore,  Mg  and  O  represent  the  number 
of  atoms  which  are  equivalent  arithmetically  to  the  ratio 
3  :  2  found  by  experiment.  Since  one  atom  of  magne- 
sium and  one  of  oxygen  unite  to  form  magnesium  oxide, 
its  formula  is  MgO.  Therefore,  the  final  equation  is  — 

Mg-f  O=MgO. 

Again,  suppose  we  wish  to  find  the  correct  equation  for 
the  reaction  between  hydrogen  and  oxygen  in  the  forma- 
tion of  water.  Experiment  shows  that  hydrogen  and  oxy- 
gen combine  in  the  ratio  of  i  :  8  by  weight.  Pursuing 
the  same  line  of  argument  as  above,  we  let  y  equal  the 
unknown  number  of  atomic  weights  of  hydrogen,  and  z 
that  of  oxygen.  The  preliminary  equation  is  — 

y  x  at.  wt.  of  hydrogen  :  z  x  at.  wt.  of  oxygen  =  i  :  8. 
The  atomic  weight  of  hydrogen  is  i  and  of  oxygen  is  16. 
The  equation  now  becomes  — 

y  x  i  :  z  x  16  =  i  :  8. 


106  Descriptive  Chemistry. 

Obviously,  y  =  2  and  z  =  i.  Now  the  symbol  H  stands 
for  i  part  of  hydrogen  and  O  for  16  parts  of  oxygen. 
Therefore,  2  H  and  O  represent  the  number  of  atoms  which 
corresponds  to  the  ratio  i  :  8,  found  by  experiment.  Since 
two  atoms  of  hydrogen  and  one  atom  of  oxygen  unite  to 
form  water,  its  formula  must  be  H2O.  Therefore  the  final 
equation  is-  H2  +  O  =  H2O. 

By  a  similar  treatment,  the  experimental  foundation  of 
all  equations  can  be  shown. 

Equations  illustrating  Reactions.  — The  simplest  equation  for  the 
preparation  of  oxygen  from  mercuric  oxide  is  — 

HgO         =      Hg     +     O  / 

Mercuric  Oxide     Mercury     Oxygen 

When  sulphur  and  carbon  are  burned  in  air  or  in  oxygen,  the  equations 

are~  s     +    o2    =         so, 

Sulphur     Oxygen     Sulphur  Dioxide 

C      +      O,      =          C02 
Carbon     Oxygen     Carbon  Dioxide 

The  equation  for  the  preparation  of  hydrogen  from  zinc  and  hydro- 
chloric acid  is  — 

Zn    +  2  HC1  =        H,         +         ZnCl,         y 

Zinc       Hydrochloric  Acid       Hydrogen       Zinc  Chloride 

When  hydrogen  burns,  the  equation  is  — 

H,        +      O       =  HaO 
Hydrogen     Oxygen     Water 

The  equation  for  the  formation  of  water  is  the  same,  though  it  is  some- 
times  written-  2  H2  +  O2  =  2  H,O. 

The  equation  for  the  reaction  in  determining  the  gravimetric  composi- 
tion of  water  is  — 

CuO        +        H2       =   H2O   +     Cu  Y 

Copper  Oxide     Hydrogen     Water     Copper 


Problems  based  on   Equations.  107 

The  interaction  of  sodium  and  water  is  represented  thus  — 

Na      +   H20  H        +         NaOH 

Sodium     Water     Hydrogen     Sodium  Hydroxide 

When  phosphorus  burns  in  air  (or  oxygen),  the  simplest  equation  is  — 

2?       +    50   -  PA 

Phosphorus     Oxygen     Phosphorus  Pentoxide 

Problems  based  on  Equations.  —  Since  equations  are  expressions 
of  chemical  reactions  which  involve  no  loss  in  weight,  it  is  possible  to 
solve  many  problems  connected  with  reactions.  An  equation  states 
the  proportions  which  participate  in  a  reaction.  Obviously,  any  con- 
venient weights  of  zinc  and  sulphuric  acid  might  be  allowed  to  interact, 
but  the  factors  and  products  are  always  in  the  proportions  given  in  the 
equation  — 

Zn  +          H2SO4       =        H2      +       ZnSO4 
Zinc     Sulphuric  Acid     Hydrogen     Zinc  Sulphate 

65  98  2  161 

This  expression  means  that  65  parts  of  zinc  always  interact  with  98 
parts  of  sulphuric  acid  and  yield  2  parts  of  hydrogen  and  161  parts  of 
zinc  sulphate.  For  parts  we  may  read  grams,  ounces,  kilograms,  —  any 
unit,  —  but  the  same  unit  must  be  used  throughout  the  calculations. 
Therefore,  if  we  know  the  weight  of  one  substance  participating  in  a 
reaction,  all  other  weights  involved  may  be  readily  calculated. 

Suppose  45  gm.  of  zinc  interact  with  sulphuric  acid ;  the  weights 
of  (a)  acid  required,  (b)  hydrogen  formed,  and  (c)  zinc  sulphate  produced 
are  found  by  the  following  proportions  :  — 

(a)  65  •:    98  {  : 45  :  y,   x=  67.8  gm.  sulphuric  acid. 
(£)  65  :      2  :  :  45  :  x,   ^=1.38  gm.  hydrogen. 
(c}  65  :  161  :  :  45  :  x,   x=  II  1.4  gm.  zinc  sulphate. 

Hence  to  solve  similar  problems,  first  write  the  equation  with  the  correct 
atomic  or  molecular  weights,1  and  then  state  the  problem  in  the  form 
of  a  proportion  like  those  given  above. 

i  The  atomic  weights  are  given  in  the  table  in  the  Appendix,  §  5.  Molecular 
weights  are  obtained  by  adding  the  proper  atomic  weights. 


io8  Descriptive  Chemistry. 

EXERCISES. 

1 .  Define  and  illustrate  the  term  equivalent.    What  is  the  equivalent 
of  hydrogen,  oxygen,  sulphur,  zinc,  copper,  magnesium,  silver,  potassium, 
aluminium  ? 

2.  What  is  the  equivalent  of  chlorine  and  of  bromine? 

3.  How  are  equivalents  determined?     Are  they  the  result  of  theory 
or  actual  analysis  ? 

4.  Expand  the  topic,  "  Atomic  weights  are  often  multiples  of  equiva- 
lents." 

5.  What  is  the  atomic  weight  of  an  element?    "How  is  it  related  to 
the  equivalent  weight  of  the  element?     Is  an  atomic  weight  absolute  or 
relative  ?     What  is  the  standard  of  atomic  weight? 

6.  What  does  O  represent  besides  one  atom  of  oxygen? 

7.  What  is  the  approximate  atomic  weight  of  hydrogen,  oxygen, 
and  sodium? 

8.  What  is  meant  by  molecular  weight?    Illustrate  by  nitric  acid  or 
potassium  chlorate.     What  is  the  relation  between  formula  and  molecular 
weight  ? 

9.  Define  and  illustrate  (a)  percentage  composition,  and  (£)  sim- 
plest formula. 

10.  How  do  we  know  that  the  correct  equation  for  the  combination 
of  magnesium  and  oxygen  is  Mg  +  O  =  MgO  ?  That  S  +  O2  =  SO2  is  the 
correct  equation  for  the  combination  of  sulphur  and  oxygen  ? 

PROBLEMS. 

1 .  Calculate  the  percentage  composition  of  (a)  water  (H2O),  (b}  zinc 
sulphide  (ZnS),  (c}  zinc   carbonate  (ZnCO3),  (d)  potassium   chlorate 
(KC103). 

2.  Calculate  the  percentage  composition  of  (#)  sugar  (C12H22O11), 
(£)  calcium  sulphate  (CaSO4),  (c)  zinc  sulphate  (ZnSO4),  (d)  magne- 
sium oxide  (MgO),  (e)  copper  oxide  (CuO). 

3.  Calculate  the  molecular  weight  of  the  following  compounds  by 
finding  the  sum  of  the  atomic  weights  :  (a)  copper  sulphate  (CuSO4), 
(If)  barium  chloride  (BaCL,),  (c)  manganese  dioxide,  (d}  calcium  oxide, 
(e)  sodium  hydroxide,  (/)  potassium  hydroxide,  (g)  sodium  carbonate, 
(^)  potassium  nitrate  (KNO3), 


Problems.  109 

4.  Calculate  the  simplest  formula  of  the  compounds  which  have  the 
indicated   composition,  and  give  the   name  of  each   compound :  (a) 
H  =  II.  11,0  =  88.89;  (£)Na  =  32.39,0  =  45-o7,S  =  22.54;  (Y)  €  =  27.27, 
0  =  72.72. 

5.  Calculate  the  simplest  formula  of  the  compounds  which  have  the 
following  composition  :  (a)  N  =  82.353,  H  =  17.647  ;  (£)  O  =  30,  Fe  =  70 ; 
(c)  H  =  i,  C  =  11.99,  O  =47-95>  K  =  39.06. 

6.  How  much  oxygen  can  be  prepared  from  (ft)  122.5  §m<  °f  potas- 
sium chlorate,  (b)  245  gm.,  and  (c)  421  gm.  ? 

Solution.     The  equation  is  — 

KC103  =  03  +  KC1 
122.5    =48  4-  74-5 

These  equation  weights  are  obtained  by  adding  the  atomic  weights  found 
in  the  table,  (a)  By  inspection,  122.5  gm-  of  potassium  chlorate  yield 
48  gm.  of  oxygen.  (b)  The  proportion  needed  is  122.5  :  48  ::  245  :  x. 
And  x  =  96  gm.  (c)  Similarly,  122.5  :  4^  :  :  421  '  *•  And  x  —  164.9. 

7.  (a)  How  much  oxygen  can  be  prepared  from  50  gm.  of  potassium 
chlorate,  and  (b)  how  much  potassium  chloride  will  remain? 

Ans.  (a}  =  19.59,  (b)  =  30.41. 

8.  A  certain  weight  of  potassium  chlorate  was  heated  until  completely 
decomposed.   The  residue  weighed  20.246  gm.    (a)  What  was  its  weight  ? 
(b)  How  much  oxygen  was  evolved?        Ans.  (a)=  33.29,  (b)  13.044. 

9.  What  weight  of  potassium  chlorate  is  needed  to  generate  144  gm. 
of  oxygen?  .^^.367.5. 

10.  What  weight  of  potassium  chloride  remains  after  obtaining  8  gm. 
of  oxygen  from  potassium  chlorate?  Am.  12.416. 

1 1 .  How  many  grams  of  oxygen  can  be  generated  from  490  gm.  of  po- 
tassium chlorate?  Ans.  192. 

12.  How  much  hydrogen  can  be  prepared  from  (a)  65  gm.  of  zinc, 
(b)  130  gm.,  (V)  297  gm.?  Ans.  (c}  9.14- 

13.  How  much  zinc  is  needed  to  prepare  (a)  2  gm.  of  hydrogen,  (b) 
14  gm.,  and  (c)  17  gm.? 

14.  How  much  zinc  sulphate  can  be  prepared  from  (a)  98  gm.  of  sul- 
phuric acid,  (b)  196  gm.,  and  (c)  427  gm.?  Ans.  (c)  701.5. 

15.  A  balloon  holds  132.74  kg.  of  hydrogen.      How  much  (a)  zinc 
and  (b)  sulphuric  acid  are  needed  to  produce  the  gas? 

Ans.  (a)  43I4-05>  (*)  6504.26. 


no  Descriptive  Chemistry. 

1 6.  How  much  (a)  mercury  and  (£)  oxygen  can  be  obtained  from 
10  gm.  of  mercuric  oxide?     (Equation  is  HgO  =  Hg  +  O,  or  216  = 
200  +  1 6.)  Ans.  (a)  9.259,  (b}  0.74. 

17.  How  much  mercury  will  remain  after  obtaining  48  gm.  of  oxygen 
by  heating  mercuric  oxide? 

1 8.  A  lump  of  carbon  weighing  24  gm.  is  burned  in  air.    What  weight 
of  (a)  carbon  dioxide  is  formed  and  (£)  oxygen  is  needed?     (c)  If  a  liter 
of  oxygen  weighs  1 .43  gm.,  what  volume  of  oxygen  is  needed  ?   (Equation 
is  C  +  O2=  CO2,  or  12  +  32  =  44.)  Ans.  (c)  44.75  1. 

19.  What  weight  of  carbon  dioxide  is  formed  by  burning  112  Ib.  of 
coal  containing  15  per  cent  of  impurities? 

20.  A  lump  of  sulphur  weighing  32  gm.  is  burned  in  air.     Calculate 
the   weight   of  (a)  oxygen   needed   and  (&)  sulphur   dioxide   formed. 
(Equation  is  S  +  O2  =  SO2,  or  32  +  32  =  64.) 

21.  Calculate  the  weight  of  oxygen  needed  to  burn  731  gin.  of  sul- 
phur containing  15  per  cent  of  impurities.  Ans.  621.35. 

22.  What  weight  of  sulphur  dioxide  is  formed  by  burning  67  per  cent 
of  8794  kg.  of  sulphur? 


CHAPTER   X. 
LIGHT,  HEAT,  ELECTRICITY,   AND  CHEMICAL  ACTION. 

CHEMICAL  action  is  always  manifested  by  one  or  more 
of  the  different  forms  of  energy,  such  as  light,  heat,  and 
electricity.  This  means  that  a  chemical  change  involves 
not  only  a  rearrangement  of  matter,  but  also  a  transfor- 
mation of  energy.  Thus,  when  coal  is  burned,  a  new 
compound  called  carbon  dioxide  is  formed,  but  heat  is  also 
liberated.  Sometimes  we  pay  more  attention  to  the  result- 
ing matter  than  to  the  energy,  but  both  are  involved.  In 
the  present  chapter  we  shall  emphasize  the  relation  of 
energy  to  chemical  action.  The  law  of  the  conservation 
of  energy  should  be  recalled  in  this  connection.  Energy, 
like  matter,  cannot  be  created  or  destroyed ;  we  can  only 
transform  it.  And  the  transformation  involves  no  loss  or 
gain.  Hence,  chemical  energy,  which  is  the  immediate 
cause  of  chemical  action,  will  appear  as  heat,  light,  or 
electricity. 

The  Relation  of  Light  to  Chemical  Action  is  illustrated  in  photogra- 
phy. Coatings  consisting  of  compounds  of  silver  and  organic  matter  are 
quickly  blackened  by  light  (see  Photography).  Sunlight  fades  many 
colors.  It  likewise  assists  the  chemical  changes  involved  in  the  growth 
of  plants.  The  formation  of  the  green  coloring  matter  of  foliage  is 
partly  due  to  sunlight.  A  mixture  of  hydrogen  and  chlorine  gases  re- 
mains unchanged  in  the  dark,  but  in  direct  sunlight  it  explodes  violently. 
On  the  other  hand,  light  is  often  a  product  of  chemical  action.  Many 
chemical  experiments  show  this,  especially  those  with  oxygen.  Sparks, 
most  flames,  and  the  flash  of  a  gun  are  other  illustrations  of  the  close 
relation  between  light  and  chemical  action.  Combustion  in  its  varied 
forms  is  also  manifested  by  light,  as  well  as  by  heat. 

HI 


H2  Descriptive  Chemistry. 


HEAT   AND    CHEMICAL    ACTION. 

Heat  and  Chemical  Action  are  closely  and  definitely 
related.  Every  chemical  change  is  attended  by  the  libera- 
tion or  absorption  of  heat.  Moreover,  the  heat  involved 
can  often  be  measured.  Heat  is  measured  in  calories,  a 
calorie  being  the  quantity  of  heat  necessary  to  raise  the 
temperature  of  I  gm.  of  water  from  o°  to  i°  C.  For 
example,  the  heat  liberated  by  the  burning  of  I  gm.  of 
hydrogen  is  34,200  cal.,  and  of  I  gm.  of  pure  charcoal  is 
about  8000  cal.  Attention  has  already  been  called  to  the 
high  temperature  of  the  hydrogen  flame  (see  Chapter  III). 

Ordinary  chemical  equations  do  not  express  changes  in  energy.  To 
represent  heat  changes,  the  number  of  calories  of  heat  involved  is 
placed  after  the  equation,  thus  :  — 

H2       +       O      =    H2O     +     68,400  cal. 
Hydrogen     Oxygen      Water 

This  is  called  a  thermal  equation,  and  it  means  that  68,400  cal.  of  heat 
are  liberated,  when  2  gm.  of  hydrogen  unite  with  16  gm.  of  oxygen  to 
form  18  gm.  of  water.  In  some  changes  heat  disappears.  Thus,  when 
carbon  unites  with  sulphur  to  form  carbon  disulphide,  heat  is  absorbed. 
The  equation  expressing  this  fact  is  — 

C     +     S2        =         CS2  19,600  cal. 

Carbon     Sulphur    Carbon  Disulphide 

Heat  involved  in  the  formation  of  a  particular  compound  is  called  heat 
of  formation  of  that  compound.  If  heat  is  liberated  in  the  formation 
of  a  compound,  the  heat  is  called  positive  (  +  );  and  the  compound  is 
termed  exothermic.  Heat  of  formation  which  is  absorbed  is  called 
negative  (  —  )  ;  and  a  compound  having  a  negative  heat  of  formation  is 
said  to  be  endothermic.  Exothermic  compounds  are  stable,  and  can 
be  decomposed  only  by  the  addition  of  the  same  quantity  of  heat  liber- 
ated by  their  formation.  Thus,  68,400  cal.  of  heat,  or  an  equivalent 
quantity  of  energy,  must  be  added  to  18  gm.  of  water  to  decompose  it 


Light,  Heat,  Electricity,  and  Chemical  Action.    113 

into  2  gm.  of  hydrogen  and  16  gm.  of  oxygen.  Such  heat  is  called 
heat  of  decomposition.  On  the  other  hand,  endothermic  compounds 
are  unstable,  and  often  explosive.  They  decompose  easily  with  the 
liberation  of  heat.  Ozone  is  endothermic.  Heat  is  absorbed  during 
its  formation  from  oxygen ;  but  when  ozone  decomposes,  heat  is  liber- 
ated. Two  parts  (by  volume)  of  ozone  form  three  parts  (by  volume) 
of  oxygen  and  liberate  72,400  cal. 

A  familiar  instance  of  the  evolution  of  heat  by  chemical 
action  is  the  slaking  of  lime.  When  lime  and  water  are 
mixed,  their  union  produces  sufficient  heat  to  boil  water 
and  often  to  set  fire  to  wood.  Steam  can  be  seen  escaping 
from  the  boxes  in  which  lime  is  being  mixed  with  water  and 
sand  to  form  plaster  or  mortar.  Buildings  in  which  lime  is 
stored  sometimes  take  fire,  if  rain  leaks  in  upon  the  lime. 
Ships  loaded  with  lime  are  in  constant  danger  of  being 
burned.  Other  substances  liberate  heat  when  added  to 
water,  e.g.  sulphuric  acid,  sodium  and  potassium  hydrox- 
ides, and  the  metals,  sodium  and  potassium. 

Heat  is  the  initial  cause  of  many  chemical  changes.  It 
is  necessary  to  start  many  reactions,  just  as  a  stone  on  top 
of  a  hill  must  be  pushed  before  it  will  roll  toward  the 
bottom.  Hydrogen  and  oxygen  mix  freely  without  com- 
bining, but  union  occurs  the  instant  heat  is  applied  in  form 
of  a  flame  or  an  electric  spark.  Similarly,  illuminating 
gas  must  be  lighted,  i.e.  raised  to  the  kindling  tempera- 
ture before  the  chemical  changes  which  cause  the  light 
and  heat  can  proceed.  These  facts  mean  that  chemical 
action  often  depends  upon  temperature.  This  statement 
has  been  strikingly  illustrated  in  the  last  four  years.  At  the 
extremely  low  temperature  obtained  by  using  liquid  air  and 
similar  substances,  it  appears  that  many  chemical  reactions 
cease.  While  at  the  exceedingly  high  temperature  pro- 
duced by  electricity  many  changes,  chemical  and  physical, 
hitherto  impossible,  occur  quickly  and  simply. 


114  Descriptive  Chemistry. 

The  Electric  Furnace  of  Moissan.  —  Until  recently  the 
heat  needed  for  chemical  changes  was  obtained  by  burn- 
ing carbon  or  its  compounds,  such  as  charcoal,  illumi- 
nating gas,  and  oil.  Sometimes  the  blast  lamp  and 
oxyhydrogen  blowpipe  were  used.  But  all  these  sources 
have  been  surpassed  in  efficiency  by  the  electric  furnace. 

It  is  well  known  that  an  electric  arc  light  produces  in- 
tense heat.  The  high  temperature  of  the  arc,  i.e.  space 
between  the  glowing  ends  of  the  carbons,  is  unequaled  by 
that  of  any  other  source  of  artificial  heat.  If  the  carbon 
rods  are  inclosed  in  a  box  that  prevents  the  escape  of  heat, 
a  temperature  estimated  to  be  about  3500°  C.  is  produced 
inside  the  box.  This  apparatus  is  called  an  electric  furnace. 
It  was  devised  and  perfected  by  the  French  chemist,  Mois- 
san, and  used  by  him  in  experimenting  at  high  temperatures. 
One  form  of  the  electric  furnace  is  shown  in  Figure  16. 


FIG.  16.  —  Moissan's  electric  furnace. 

Moissan's  description  of  this  furnace  is  as  follows :  "  It  consisted  of 
two  bricks  of  quicklime  placed  one  on  top  of  the  other.  The  lower 
brick  contained  a  longitudinal  groove  to  receive  the  two  electrodes 
[carbon  rods],  and  situated  in  the  center  was  a  small  cavity.  This 
cavity  might  vary  in  size,  and  contained  a  bed  some  centimeters  in 
depth  of  the  substance  to  be  acted  upon  by  the  heat  of  the  arc,  or  a 
small  crucible  of  carbon  containing  the  substance  to  be  treated  may  be 
placed  there.  The  upper  brick  was  slightly  hollowed  out  in  the  part 
just  above  the  arc.  As  the  intense  heat  of  the  current  soon  melted  the 


HENRI    MOISSAN 

1852 

THE    EMINENT   FRENCH    CHEMIST    WHOSE    DISCOVERIES    CONTINUE   TO    ENRICH 
INORGANIC    CHEMISTRY 


Light,  Heat,  Electricity,  and  Chemical  Action.    115 


surface  of  the  lime,  giving  it,  at  the  same  time,  a  beautiful  polish,  a 
dome  was  obtained  in  this  way  which  reflected  all  the  heat  on  to  the 
small  cavity  which  contained  the  crucible."  Figure  17  is  a  vertical  sec- 
tion of  the  furnace,  showing  the  parts  slightly  separated.  The  furnace 
is  small,  some  being  only  16  to  18  cm.  (about  7  in.)  long,  15  cm.  wide, 
and  14  cm.  high. 
The  carbon  rods  are 
from  i  to  5  cm.  in 
diameter. 

When    a    cur- 
rent    is     passed 

through    the   Car-     FIG.  17.  —  Vertical  section  of  Moissan's  electric  furnace. 

bon      rods,      the 

tremendous  heat  produced  is  retained  in  the  space  by  the 
non-conducting  walls  and  acts  upon  the  substance  below 
the  arc.  The  outside  of  the  furnace  remains  cold  enough 
to  be  touched  by  the  hand,  but  the  inside  is  almost  twice 
as  hot  as  the  oxyhydrogen  flame.  There  is  no  electrical 
action  upon  the  chemicals.  The  intense  heat  alone  pro- 
duces the  remarkable  changes,  which  are  often  accom- 
plished in  a  few  minutes.  Sand,  lime,  magnesium  oxide, 
and  other  refractory  oxides  melt  and  volatilize.  The  ele- 
ments carbon,  silicon,  and  boron  boil;  and  gold,  copper, 
and  platinum  quickly  melt  and  vaporize.  Large  masses 
of  rare  and  uncommon  elements  are  quickly  reduced  from 
their  oxides  and  obtained  in  the  pure  state,  e.g.  chromium, 
manganese,  tungsten,  uranium,  and  molybdenum.  Char- 
coal becomes  graphite.  And  stable  compounds  of  carbon, 
boron,  and  silicon  are  formed.  These  are  the  carbides, 
borides,  and  silicides.  Some  of  the  carbides  have  an  in- 
dustrial use  as  well  as  scientific  interest,  especially  calcium 
carbide  and  silicon  carbide  (see  below).  Other  carbides 
are  the  sources  of  pure  metals,  since  the  fusion  of  a  car- 
bide and  oxide  of  the  same  metal  yields  the  metal  itself. 


ii6  Descriptive  Chemistry. 

Industrial  Use  of  the  Electric  Furnace.  —  Huge  elec- 
tric furnaces  constructed  on  the  type  devised  by  Moissan 
are  in  active  operation.  And  since  electricity  is  now  ob- 
tained in  many  localities  by  running  dynamos  by  water, 
new  industries  requiring  intense  and  continuous  heat  have 
recently  sprung  into  existence.  Several  of  these  plants 
are  located  at  Niagara  Falls,  which  furnishes  enormous 
power  at  a  relatively  small  expense. 

Calcium  Carbide  is  made  on  a  large  scale  by  heating  a 
mixture  of  lime  and  coke  (a  form  of  carbon)  in  an  electric 
furnace.  The  chemical  change  is  caused  solely  by  the 
intense  heat  and  may  be  represented  thus  :  — 

3C       +       CaO       =         CaC2          +          CO 
Carbon  Lime  Calcium  Carbide       Carbon  Monoxide 

This  method  of  making  calcium  carbide  cheaply  was  dis- 
covered independently  and  at  about  the  same  time  (1892- 
1895)  by  Moissan  and  Willson.  The  furnaces  now  in 
operation  vary  in  details,  but  all  have  one  essential  feature, 
viz.,  the  heat  is  generated  by  an  electric  current  passing 
between  two  carbon  electrodes.  In  most  furnaces  one  elec- 
trode is  a  crucible  wholly  or  partly  of  carbon,  and  the  other 
electrode  is  a  stout  carbon  pillar  dipping  into  the  mixture. 
Calcium  carbide  is  a  hard,  brittle,  dark  gray,  crystalline 
solid  with  a  metallic  luster.  Its  specific  gravity  is  2.2. 
The  most  striking  and  useful  property  is  its  action  with 
water,  acetylene  being  formed,  thus:  — 

CaC2      +     2H20     =    C2H2          -f         Ca(OH)2 
Calcium  Carbide        Water  Acetylene  Calcium  Hydroxide 

Calcium  carbide  is  used  to  generate  acetylene  gas.  This 
gas  burns  with  a  brilliant  flame,  and  is  coming  into  general 
use  as  an  illuminant.  Owing  to  its  action  with  water, 


Light,  Heat,  Electricity,  and  Chemical  Action.    117 

calcium  carbide  is  packed  and  sold  in  air-tight  cans  (see 
Acetylene). 

Carborundum  is  a  compound  of  silicon  and  carbon,  hav- 
ing the  composition  SiC.  It  is  made  in  the  electric  furnace 
by  fusing  sand  (silicon  dioxide,  SiO2),  coke,  saw.dust,  and 
common  salt.  The  essential  chemical  change  is  repre- 
sented thus :  — 

SiO2         +       3C       =       SiC  +          2  CO 

Silicon  Dioxide  Carbon  Carborundum  Carbon  Monoxide 
Carborundum  is  silicon  carbide  (or  carbon  silicide).  It 
is  a  crystallized  solid,  varying  in  color  from  white  to  emerald 
green  and  is  sometimes  iridescent.  It  is  extremely  hard, 
being  harder  than  ruby  and  nearly  as  hard  as  diamond. 
Hence  it  is  made  into  grinding  wheels,  whetstones,  and 
polishing  cloths.  Over  three  million  pounds  were  made  at 
Niagara  Falls  in  1902,  and  the  output  is  constantly 
increasing. 

Carborundum  is  a  good  conductor  of  heat.  Its  specific  gravity  is 
about  three.  Acids  have  no  action  upon  it,  but  it  is  decomposed  by 
fusing  with  potassium  hydroxide  and  other  alkalies. 

Carborundum  is  manufactured  in  a  huge  electric  furnace,  shown  in 
Figure  18.  It  is  an  oblong  box  of  bricks  with  permanent  ends  and  loosely 
built  sides.  Each  end  is  provided  with  a  heavy  metal  plate.  The  wires 
for  the  electric  current  are  attached  to  the  outer  ends  of  these  plates, 
while  the  huge  carbon  electrodes  fit  into  the  inner  ends,  and  project  into 
the  furnace.  A  cylinder  of  granulated  coke  makes  an  electrical  connec- 
tion between  the  electrodes.  In  this  furnace  the  rnixture  is  not  heated 
by  an  electrical  arc,  but  by  the  resistance  of  the  carbon  core  to  the  pas- 
sage of  the  powerful  current  of  electricity.  The  chemical  change,  as  in 
the  manufacture  of  calcium  carbide,  is  due  solely  to  heat.  The  current 
is  passed  through  the  mixture  for  about  eight  hours.  When  the  opera- 
tion is  over  and  the  furnace  is  cool,  the  side  walls  are  pulled  down,  and 
the  carborundum  is  removed.  The  purest  grade  is  found  around  the 
core.  It  is  crushed,  treated  with  sulphuric  acid  to  remove  the  impurities, 
washed,  dried,  and  graded  according  to  the  size  of  the  particles. 


n8 


Descriptive  Chemistry. 


Artificial  Graphite  is  formed  in  the  manufacture  of 
carborundum.  It  is  also  made  by  heating  a  certain  grade 
of  anthracite  coal  in  an  electric  furnace.  It  is  extensively 
used  in  making  electrodes  for  electric  furnaces.  Over 
800,000  Ib.  were  manufactured  in  1902  at  Niagara  Falls. 
Graphite  is  a  form  of  carbon  (see  Graphite). 


Light,  Heat,  Electricity,  and  Chemical  Action.    119 


ELECTRICITY    AND    CHEMICAL    ACTION. 

The  Relation  between  Electricity  and  Chemical  Action 

has  always  been  a  fascinating  subject.  Volta  constructed 
his  voltaic  pile  about  1800.  This  was  one  of  the  first,  per- 
haps the  first,  source  of  an  electric  current.  In  May,  1800, 
Nicholson  and  Carlisle  decomposed  water  into  hydrogen 
and  oxygen  by  an  electric  current  obtained  from  a  thermo- 
pile. In  the  same  year  Cruikshank  obtained  lead  and 
copper  from  solutions  of  their  salts.  And  in  1807  Davy 
isolated  the  elements,  sodium  and  potassium,  by  passing  an 
electric  current  (obtained  from  a  large  battery)  through 
fused  caustic  soda  and  caustic  potash  respectively.  From 
that  time  until  the  present  day,  the  relation  between  elec- 
tricity and  chemical  action  has  engaged  the  attention  of 
chemists.  And  their  labors  have  built  up  a  branch  of 
chemistry  called  electrochemistry,  which 
has  recently  attained  considerable  com- 
mercial importance. 

The  Voltaic  (or  Galvanic)  Cell  in  its  simplest 
form  consists  of  two  metals  connected  by  a  wire 
and  dipped  into  a  liquid  which  will  interact  with 
one  of  the  metals  (Fig.  19).  Copper,  zinc,  and 
water  containing  sulphuric  acid  may  be  used  as 
an  illustration.  When  the  connected  metals  are  FlG  I9._voltaic  cell, 
put  into  the  acid,  the  zinc  slowly  disappears  and 

hydrogen  bubbles  appear  on  the  copper.  Further  examination  would 
show  that  the  zinc  and  sulphuric  acid  interacted,  forming  zinc  sulphate. 
The  chemical  change  is  the  one  already  described  under  hydrogen,  and 
may  be  represented  thus :  — 

Zn  +  H2S04  H2  +  ZnS04 

Zinc  Sulphuric  Acid  Hydrogen  Zinc  Sulphate 

The  connecting  wire  becomes  electrified  and  exhibits  the  effects  of  an 
electric  current,  viz.,  it  becomes  warm,  it  makes  a  magnetic  needle  move, 


I2O  Descriptive  Chemistry. 

and  a  shower  of  sparks  is  produced  if  the  wire  is  cut  and  one  end  is 
drawn  down  a  file  while  the  other  is  held  firmly  upon  it.  The  source  of 
the  electric  current  is  obviously  the  chemical  action  between  the  acid  and 
zinc.  The  copper  is  necessary,  otherwise  the  product  of  the  chemical 
action  would  be  merely  heat.  Carbon  is  often  used  in  place  of  copper, 
and  other  liquids  instead  of  sulphuric  acid.  The  liquid  chosen,  how- 
ever, must  be  one  that  will  interact  with  zinc  or  its  substitute.  Several 
cells  joined  together  form  an  electric  battery.  For  many  years  the 
battery  was  the  chief  source  of  the  electric  current.  And  it  is  now  used, 
especially  for  ringing  telephone,  house,  fire  alarm,  and  signal  bells,  and 
in  operating  the  telegraph.  The  dynamo  is  now  widely  used  to  generate 
powerful  currents  of  electricity. 

Electrochemical  Terms.  —  Faraday  (1791-1867)  investi- 
gated electrochemistry  about  1834,  and  introduced  many 
terms  in  common  use.  He  called  the  decomposing  process 
electrolysis,  and  the  decomposable  liquid  the  electrolyte ;  the 
wire  by  which  the  current  entered  he  called  the  anode ; 
and  that  by  which  it  escaped,  the  cathode.  "  Finally,"  he 
says,  ^require  a  term  to  express  those  bodies  which  pass 
to  the  electrodes)  I  propose  to  distinguish  such  bodies  by 
calling  those  anions  which  go  to  the  anode  of  the  decom- 
posing body ;  and  those  passing  to  the  cathode,  cations ; 
and  when  I  have  occasion  to  speak  of  these  together,  I 
shall  call  them  ions.  Thus,  chloride  of  lead  is  an  electro- 
lyte, and  when  electrolyzed  evolves  the  two  ions,  chlorine 
and  lead,  the  former  being  an  anion  and  the  latter  a 
cation."  These  terms  are  so  used  to-day,  but  they  demand 
a  broader  definition.  Electrolysis  is  the  series  of  chemi- 
cal changes  caused  by  the  passage  of  an  electric  current 
through  a  dissolved  or  fused  (i.e.  melted)  compound.  The 
compound  thus  decomposed  is  an  electrolyte.  The  metallic 
or  carbon  rods  which  conduct  the  current  of  electricity  to 
and  from  the  electrolyte  are  called  the  poles,  or  better,  the 
electrodes.  Electrodes  are  usually  made  of  platinum,  cop- 


Light,  Heat,  Electricity,  and  Chemical  Action.    121 

per,  zinc,  mercury,  or  hardened  carbon ;  they  may  have 
any  shape  —  rod,  wire,  sheet,  plate,  box,  crucible  ;  and  they 
may  also  be  solid,  liquid,  or  powder,  as  well  as  fixed  or 
movable.  The  electrodes  are  connected  by  wires  with  the 
source  of  the  electric  current,  and  serve  as  "doors"  —  to 
quote  Faraday  again  —  for  the  current  to  flow  into  and  out 
of  the  electrolyte  and  through  the  wire  connecting  the 
electrodes.  We  speak  of  a  " current"  of  electricity  and 
of  electricity  as  "  flowing,"  although  we  do  not  know  the 
nature  of  electricity,  nor  do  we  mean  really  that  it  flows,  like 
a  river,  only  in  one  direction.  It  is  customary  to  speak  of 
the  current  as  entering  the  electrolyte  by  the  anode  or 
positive  electrode  and  leaving  by  the  negative  electrode 
or  cathode.  The  anode  is  the  electrode  that  is  often  con- 
sumed or  worn  away,  either  mechanically  or  chemically. 
But  solids  are  often  deposited  upon  the  cathode,  as  will 
soon  be  described,  ^ons  are  those  parts  of  the  decomposed 
electrolyte  which  are  believed  to  be  material  carriers  of  elec- 
tricij^y Aw  comes  from  a  Greek  word  which  means  wander- 
ing or  migrating.  And  a  cation  is  that  ion  which  moves 
down  or  along  with  the  current  of  electricity  to  the  cathode 
where  it  is  separated,  deposited,  or  modified;  while  an  anion 
is  that  ion  which  moves  upward  or  against  the  current  to  the 
anode,  where  it  likewise  appears  in  various  forms.  Anions 
are  electro-negative  ions,  but  cations  are  electro-positive 
ions.  Metallic  ions  are  cations  ;  hence  metals  are  deposited 
at  the  negative  electrode  or  cathode.  Non-metallic  ions 
are  usually  anions,  therefore  oxygen,  chlorine,  and  their 
oxides  and  hydroxides  appear  at  the  anode.  Hydrogen  is 
electro-positive.  In  general,  metals  are  electro-positive, 
and  non-metals  (except  hydrogen)  are  electro-negative. 
Ions  follow  the  law  of  electric  attraction  and  repulsion, 
viz.,  ions  with  the  same  kind  of  electrification  repel  each 


122 


Descriptive  Chemistry. 


FIG.  20. —  Electrolytic  cell.  A 
and  C  are  the  electrodes,  R  is  the 
electrolyte,  B  or  D  is  the  battery 
or  dynamo. 


other,  and  those  with  unlike  kinds  attract.  Hence  the 
electro-positive  cations  move  toward  the  electro-negative 
cathode,  and  the  electro-negative  anions  move  toward  the 
electro-positive  anode.  Ions  are  further  described  under 

lonization  (see  below).  An  elec- 
trolytic cell  is  the  apparatus  in 
which  electrolysis  takes  place 
(Fig.  20).  Its  parts  are  analo- 
gous to  the  voltaic  cell.  There 
must  be  a  containing  Vessel,  the 
two  electrodes,  and  the  electro- 
lyte. The  vessel  may  have  any 
desired  shape,  and  is  made  of 
material  which  will  resist  the  corrosive  action  of  the  electro- 
lyte or  which  will  withstand  a  high  temperature.  Unlike  the 
voltaic  cell,  the  electrolytic  cell  generates  no  electric  current ; 
it  receives  the  current  from  a  dynamo  or  a  battery.  Elec- 
trolysis is  accomplished  on  a  large  scale  in  electrolytic  cells. 

Illustrations  of  Electrolysis.  —  Electrolysis  may  be 
simple,  but  it  is  usually  very  complex.  Two  illustrations 
will  be  given.  When  two  platinum  electrodes  are  put  into 
melted  zinc  chloride  and  a  current  of  electricity  is  passed, 
zinc  is  deposited  at  the  cathode,  and  chlorine  gas  is  liber- 
ated at  the  anode.  This  is  a  simple  instance  of  electrolysis. 
But  when  an  aqueous  solution  of  sodium  chloride  is  electro- 
lyzed,  the  action  is  different.  Theoretically,  the  products 
should  be  sodium  and  chlorine,  but  they  are  hydrogen, 
sodium  hydroxide,  and  chlorine.  The  sodium  separated  at 
the  cathode  immediately  interacts  with  the  water  to  form 
hydrogen  and  sodium  hydroxide.  Furthermore,  unless  the 
chlorine  and  sodium  hydroxide  are  removed,  they  will 
interact  to  form  compounds  of  chlorine,  which  vary  in 
composition  with  the  temperature,  etc. 


Light,  Heat,  Electricity,  and  Chemical  Action.    123 

The  Electrolysis  of  Water  is  more  complex  than  is  ordinarily  sup- 
posed. Strictly  speaking,  it  is  the  sulphuric  acid,  and  not  the  water, 
that  is  electrolyzed.  Perfectly  pure  water  does  not  conduct  electricity, 
and  is  consequently  not  decomposed  by  it.  But  since  the  same  amount 
of  sulphuric  acid  is  always  present,  no  matter  how  long  the  action  con- 
tinues, it  is  customary  to  speak  of  the  total  change  as  the  electrolysis  of 
water.  The  hydrogen  and  oxygen  gases,  which  collect  at  the  cathode 
and  anode  respectively,  are  merely  the  end  products  of  a  series  of 
changes.  Small  quantities  of  ozone  and  hydrogen  dioxide  are  also 
formed. 

Faraday's  Law.  —  In  his  study  of  electrolysis,  Faraday  found  that  a 
measured  quantity  of  electricity  liberated  different  but  definite  amounts 
of  the  chemical  elements.  For  example,  the  current  which  liberated 
i  gm.  of  hydrogen  also  liberated  8  gm.  of  oxygen,  35.5  gm.  of  chlorine, 
108  gm.  of  silver,  31.7  gm.  of  copper,  and  so  on.  These  numbers  are 
identical  with  the  chemical  equivalents  of  these  .  elements  (compare 
Equivalents,  Chapter  IX) .  Faraday  called  them  electrochemical  equiv- 
alents, to  emphasize  their  chemical  and  electrical  relationship.  But  the 
term  electrochemical  equivalent  now  means,  however,  the  weight  of  an 
element  deposited  or  liberated  by  a  current  of  a  certain  arbitrary  value 
(i  ampere  in  I  second) .  For  example,  the  electrochemical  equivalent  of 
hydrogen  is  0.000010441  gm.,  of  oxygen  is  0.00008287,  and  sometimes 
0.00016574,  of  copper  is  0.0003294,  and  sometimes  0.0006588,  of  silver  is 
0.001118.  This  general  relation  is  often  stated  as  Faraday's  Law, 
thus :  — 

When  the  same  quantity  of  electricity  acts  upon  different  electrolytes, 
the  ratio  between  the  quantities  of  liberated  products  is  the  same  as 
between  their  chemical  equivalents. 

Faraday  also  showed  that  the  amount  of  decomposition  —  the  chem- 
ical work,  we  might  say  —  is  proportional  to  the  total  amount  of  elec- 
tricity used.  It  makes  no  difference  whether  the  current  is  strong  or 
weak,  nor  whether  the  time  of  its  flow  is  long  or  short.  A  certain 
quantity  of  electricity  will  do  so  much  chemical  work  —  no  more  and 
no  less.  Thus  a  given  quantity  of  electricity  passed  through  copper 
sulphate  solution  always  deposits  the  same  weight  of  copper  at  the 
cathode.  These  two  principles  of  Faraday  are  at  the  foundation  of  all 
electrochemical  industries.  Their  importance  can  hardly  be  over- 
estimated. 


124  Descriptive  Chemistry. 

Industrial  Applications  of  Electrolysis.  — The  earliest 
industrial  application  of  electrolysis  was  in  electrotyping 
and  electroplating.  These  operations  consist  in  depositing 
a  thin  film  of  metal  upon  a  surface.  They  are  fundamen- 
tally the  same,  though  copper  is  the  only  metal  used  for 
producing  electrotypes.  Electrotypes  are  exact  repro- 
ductions of  the  original  objects.  The  process  of  electro- 
typing  is  substantially  as  follows  :  the  page  of  type,  or 
the  woodcut,  is  first  reproduced  in  wax  or  plaster.  This 
exact  impression  is  next  covered  with  powdered  graphite 
to  make  it  conduct  electricity.  The  coated  mold  is  then 
suspended  as  the  cathode  in  an  acid  solution  of  copper 
sulphate ;  the  anode  is  a  plate  or  bar  of  copper.  When 
the  current  is  passed,  electrolysis  occurs ;  copper  is  dis- 
solved from  the  anode  and  deposited  upon  the  mold  in  a 
film  of  any  desired  thickness.  The  exact  copper  copy  is 
stripped  from  the  mold,  backed  with  metal  and  mounted  on 
a  wooden  block,  and  used  instead  of  the  type  or  woodcut 
itself.  By  this  process  exact  copies  of  expensive  wood 
engravings  can  be  cheaply  reproduced,  and  type  can  be 
saved  from  the  wear  and  tear  of  printing.  Most  books, 
magazines,  and  newspapers  are  now  printed  from  electro- 
types. The  process  of  electroplating  differs  from  elec- 
trotyping in  only  one  essential,  viz.,  in  electroplating,  the 
deposited  film  is  not  removed  from  the  object.  The  object 
to  be  plated  is  carefully  cleaned  and  made  the  cathode ; 
the  anode  is  a  bar  or  plate  of  the  metal  to  be  deposited. 
When  the  current  passes  through  the  system,  the  metal  is 
firmly  deposited  upon  the  object.  The  electrolysis  would 
take  place,  of  course,  if  any  anode  were  present ;  but  anodes 
of  the  metal  to  be  deposited  are  usually  used  to  prevent 
the  solution  or  "  bath  "  from  weakening.  They  accom- 
plish the  purpose  by  replenishing  the  solution  with  metal 


Light,  Heat,  Electricity,  and  Chemical  Action.    125 

as  fast  as  it  is  removed  and  deposited  upon  the  cathode. 
Silver,  nickel,  and  gold  are  the  usual  metals  used  in 
electroplating  (see  these  metals). 

Electroplating  and  electrotyping  have  been  done  since 
about  1840.  It  is  only  within  the  last  ten  or  fifteen  years, 
however,  that  the  electric  current  has  been  profitably 
applied  in  many  industries.  But  during  this  time  the 
development  of  electrochemistry  has  been  very  marked. 
The  largest  of  these  industries  is  the  refining  of  copper. 
The  process  is  similar  to  that  described  under  electro- 
typing.  Other  metals,  such  as  gold,  silver,  and  lead,  are 
extracted  from  their  ores  and  purified  by  electricity,  though 
the  older  processes  are  still  used.  All  the  aluminium,  mag- 
nesium, and  sodium  of  commerce  are  now  manufactured  by 
passing  an  electric  current  through  their  fused  compounds. 
Nearly  all  the  domestic  potassium  chlorate  and  much  of 
the  caustic  soda  are  made  by  electricity.  The  same  is 
true  of  barium  compounds  and  many  other  chemicals. 
These  electrochemical  processes  will  be  fully  discussed  in 
the  appropriate  places. 

The  Theory  of  Electrolysis.  —  Many  theories  have  been 
proposed  to  explain  electrolysis.  According  to  the  theory 
now  generally  held,  electrolysis  is  not  the  splitting  or  tear- 
ing apart  of  molecules  by  the  electric  current.  It  is  the 
carrying  of  electricity  from  one  electrode  to  the  other  by 
ions.  Dissolved  or  fused  compounds  are  more  or  less  dis- 
sociated into  ions  before  the  current  of  electricity  is  intro- 
duced,/and  the  current  flows  simply  because  the  ions  are 
there  to  carry  it.  Since  these  ions  are  charged  with  elec- 
tricity, the  dissociation  is  called  electrolytic  dissociation 
or  ionization.  Ions  are  not  atoms,  but  electrically  charged 
atoms  or  groups  of  atoms.  Thus,  when  sodium  chloride  is 
dissolved  in  water,  much  of  the  salt  dissociates  into  the 


126  Descriptive  Chemistry. 

ions,  sodium  and  chlorine ;  the  sodium  ions  are  charged 
positively,  and  the  chlorine  atoms  negatively.  Now,  when 
an  electric  current  is  passed  into  the  solution,  the  ions 
move  toward  their  proper  electrodes,  carrying  the  electric 
charges  with  them.  In  brief,  the  current  sorts  the  ions, 
which  in  turn  migrate  with  their  charges.  When  the  ions 
reach  their  respective  electrodes,  they  give  up  their  electric 
charges  and  assume  their  normal  conditions.  Thus,  the 
positive  sodium  ions  give  up  their  charges  at  the  negative 
electrode,  or  cathode,  and  become  sodium  atoms.  The 
latter  interact  with  water  to  form  hydrogen  and  sodium 
hydroxide.  Similarly,  the  negative  chlorine  ions  give  up 
their  charges  at  the  positive  electrode,  or  anode,  and 
become  neutral  atoms,  which  at  once  unite  to  form  chlo- 
rine molecules. 

Electrolysis  and  Solution.  —  According  to  the  above 
theory,  the  properties  of  many  water  solutions  are  closely 
related  to  the  phenomena  of  electrolysis.  For  many  years 
it  was  believed  that  a  dissolved  substance  was  distributed 
unchanged  throughout  the  solvent.  It  was  also  believed 
that  certain  dissolved  substances  combined  in  part  with 
the  water  —  a  view  held  to-day.  The  first  real  step  toward 
a  settlement  of  the  problem  was  taken  when  the  electri- 
cal conductivity  of  solutions  was  compared.  Experiments 
show  that  the  electrical  conductivity  of  solutions  varies 
between  wide  limits.  Water  itself  is  practically  a  non- 
conductor, a  sugar  solution  is  a  very  poor  conductor,  while 
solutions  of  most  acids,  bases,  and  salts  are  excellent  con- 
ductors. Water  solutions,  therefore,  are  of  two  kinds : 
(i)  those  which  conduct  electricity,  and  (2)  those  which 
do  not,  or  only  very  slightly.  But  we  have  already  seen 
that  the  first  class  consists  of  electrolytes.  Hence,  two 
things  are  believed  about  water  solutions:  (i)  that  when 


Light,  Heat,  Electricity,  and  Chemical  Action.    127 

acids,  bases,  and  salts  are  dissolved  in  water,  they  are  dis- 
sociated into  ions,  and  (2)  that  when  sugar  and  similar 
substances  are  dissolved  in  water  they  dissociate  very 
slightly  or  not  at  all.  The  amount  of  dissociation  depends 
largely  upon  the  relative  amounts  of  solute  and  solvent, 
i.e.  upon  the  dilution  of  the  solution.  The  dissociation  is 
slight  in  concentrated  solutions,  but  increases  as  the  dilu- 
tion increases.  Not  all  acids,  bases,  and  salts  dissociate  to 
the  same  degree.  The  percentage  of  dissociation  of  some 
of  these  compounds  in  solutions  of  a  certain  strength  and  at 
the  same  temperature  (i 8°  C.)  is  given  in  the  following  — 

TABLE  OF  IGNIZATION. 


SUBSTANCE.  . 

PER  CENT  OF  IONIZATION. 

Hydrochloric  acid 

78 

/° 
82 

Potassium  chloride    

7C 

Potassium  nitrate       

/  j 
64. 

Potassium  hydroxide 

77 

Sodium  hydroxide 

7? 

Numerous  facts  support  the  theory  of  ionization.  (i) 
Varying  electrical  conductivity  has  already  been  mentioned. 
(2)  It  has  long  been  known  that  solutions  boil  at  a  higher 
temperature  and  freeze  at  a  lower  temperature  than  pure 
water.  A  fresh-water  river,  for  example,  freezes  before 
the  ocean,  and  water  containing  considerable  mineral  mat- 
ter boils  at  a  higher  temperature  than  pure  drinking  water. 
It  is  generally  true  that  a  dissolved  substance  raises  the 
boiling  point  and  loivers  the  freezing  point  of  a  given  solu- 
tion. Now,  when  weights  of  substances  proportional  to 


128  Descriptive  Chemistry. 

their  molecular  weights  are  dissolved  in  the  same  volume  of 
water,  the  boiling  point  of  each  solution  is  raised  the  same 
number  of  degrees  and  the  freezing  point  is  lowered  the 
same  number  of  degrees.  These  facts  are  now  applied 
experimentally  to  determine  molecular  weights.  In  many 
cases  the  molecular  weights  thus  found  agree  with  the 
values  obtained  by  other  methods.  Thus,  if  X  is  the 
depression  produced  by  a  one  per  cent  solution  of  sugar, 
and  Y  the  depression  produced  by  a  one  per  cent  solution 
of  urea,  the  following  proportion  may  be  written,  because 
the  depressions  of  the  freezing  points  are  inversely  propor- 
tional to  the  molecular  weights  — 

Y:  X  :  :  mol.  wt.  of  sugar :  mol.  wt.  of  urea. 

The  molecular  weight  of  sugar  is  known  to  be  342,  and 
from  the  proportion  the  molecular  weight  of  urea  is  60, 
which  agrees  with  that  found  by  other  methods.  This 
method  is  applicable  to  many  compounds  and  is  helpful 
in  deciding  whether  a  molecular  weight  is  a  given  number 
or  its  multiple.  There  is  a  marked  disagreement  to  this 
rule,  however,  in  the  case  of  solutions  of  acids,  bases,  and 
salts.  That  is,  electrolytes  are  exceptions.  In  some 
instances  the  molecular  weight  is  only  half  that  found 
by  other  methods.  Thus,  the  molecular  weight  of  sodium 
chloride  was  found  to  be  about  30,  instead  of  58.5 — the 
correct  molecular  weight.  Hence,  it  is  believed  that  the 
solutions  of  acids,  bases,  and  salts  contain  ions  which  act 
like  molecules  in  their  effect  upon  the  freezing  and  boiling 
points  of  solutions.  The  behavior  of  acids,  bases,  and 
salts  in  solution  led  the  Swedish  chemist  Arrhenius,  in  1887, 
to  extend  the  ideas  of  Faraday  and  to  propose  the  present 
theory  of  solution. 


Light,  Heat,  Electricity,  and  Chemical  Action.    129 

Application  of  the  Theory  of  lonization.  —  Many  ob- 
scure facts  of  chemistry  become  intelligible  when  inter- 
preted by  the  theory  of  ionization.  (i)  Ordinary  tests  are 
tests  for  ions.  For  example,  all  chlorides  in  solution  have 
the  same  test.  That  is,  they  all  interact  with  silver  nitrate 
in  solution,  because  all  have  chlorine  ions  in  the  solution. 
Similarly,  all  soluble  sulphates  interact  with  barium  chlo- 
ride in  solution,  because  all  sulphates  have  SO4  ions  in  the 
solution.  Both  silver  chloride  and  barium  sulphate  are 
insoluble,  and  are  removed  from  the  solution  as  precipi- 
tates. A  complete  illustration  will  make  this  fact  clearer. 
The  silver  nitrate  and  sodium  chloride  solutions  before 
mixing  consist  largely  of  the  ions  of  silver,  NO3-group, 
sodium,  and  chlorine.  When  mixed,  the  ions  of  silver 
and  chlorine  unite  to  form  silver  chloride,  which  is  in- 
soluble and  hence  not  ionized ;  the  solution  still  contains 
ions  of  sodium  and  of  the  NO3-group.  On  the  other 
hand,  if  solutions  of  potassium  chlorate  and  silver  nitrate 
are  mixed,  no  silver  chloride  is  formed,  because  no  chlo- 
rine ions  are  available.  Potassium  chlorate  dissociates 
into  ions  of  potassium  and  C1O3.  Equations  are  often 
used  to  express  ionization.  Thus,  the  ionic  equation  for 
the  interaction  of  sodium  chloride  and  silver  nitrate  is  — 

Na  +  Cl  4  Ag  +  NO8  =  AgCl  4-  Na  4  NO3. 

(2)  lonization  explains  the  General  Properties  of  Acids, 

Bases,  and  Salts.     Acids  in  solution  turn  litmus  red,  be- 

+ 

cause  their  solutions  contain  hydrogen  ions  (H).  Simi- 
larly, bases  turn  litmus  blue,  because  their  solutions  contain 

hydroxyl  ions  (OH).  But  solutions  of  neutral  salts  con- 
tain neither  hydrogen  nor  hydroxyl  ions,  hence  they  do 
not  affect  litmus.  The  above  principles  can  be  readily 


i jo  Descriptive  Chemistry. 

extended  to  cover  acid  and  basic  salts.  The  other  general 
properties  of  acids  and  bases  are  believed  to  be  due  to  the 
above  causes.  (3)  Neutralization,  interpreted  by  the  ionic 
theory,  is  fundamentally  the  union  of  hydrogen  and  hy- 
droxyl  ions  to  form  molecules  of  water.  Suppose  hydro- 
chloric acid  and  potassium  hydroxide  are  mixed.  The 
solution  at  first  contains  the  hydrogen,  chlorine,  potassium, 
and  hydroxyl  —  all  as  ions.  But  the  hydrogen  and  hy- 
droxyl  immediately  unite  to  form  water,  leaving  the  po- 
tassium and  chlorine  ions  in  the  solution.  This  solution 
is  thus  rendered  neutral  by  the  removal  of  the  hydrogen 
ion  —  its  acid  constituent  —  and  of  the  hydroxyl  ion  —  its 
basic  constituent.  The  ionic  equation  expressing  the 
neutralization  of  potassium  hydroxide  by  hydrochloric 
acid  is  — 

K  +  OH  +  H  +  Cl  =  K  +  Cl  +  H2O. 
The    potassium    and    chlorine    ions   remain  free  and    un- 
combined  until  the  solution  is  evaporated.     As  the  con- 
centration increases,  the  ions  unite  until  nothing  remains 
except*  the  neutral  salt  potassium  chloride. 

Neutralization,  therefore,  as  interpreted  by  the  ionic  theory,  is  essen- 
tially a  union  of  hydroxyl  and  hydrogen  ions.  This  view  is  supported 
by  much  experimental  evidence.  For  example,  the  heat  of  neutraliza- 
tion produced  by  the  interaction  of  equivalent  quantities  of  strong 
acids  and  bases  is  approximately  the  same. 

EXERCISES. 

1.  What  transformations    of  energy  accompany  chemical   action? 
Illustrate  your  answer. 

2.  State  and  illustrate  the  law  of  the  conservation  of  energy. 

3.  Discuss  the  relation  of  light  to  chemical  action.     Give  popular 
and  scientific  illustrations  of  (a}  the  production  of  chemical  action  by 
light,  and  (#)  production  of  light  by  chemical  action. 

4.  Define  and  illustrate  (a}  calorie,  (6)  thermal  equation,  (c}  heat 
of  formation,  (d}  exothermic,  (e)  heat  of  decomposition,  (/)  endothermic. 


Light,  Heat,  Electricity,  and  Chemical  Action.    131 

5.  Give  several  illustrations  of  the  production  of  (a}  heat  by  chemi- 
cal action,  and  (b)  vice  versa. 

6.  When  an  electric  spark  is  passed  through  a  mixture  of  two  vol- 
umes of  hydrogen  and  one  volume  of  oxygen,  what  is  the  result?     Is  it 
due  directly  to  electricity  or  to  heat? 

7.  Define  and  illustrate  kindling  temperature. 

8.  Name  several  sources  of  heat.     How  may  electricity  be  used  as 
a  source  of  heat  ? 

9.  Describe  Moissan's  electric  furnace.     Why  is  it  so  efficient?     Is 
its  effect  thermal  or  electrical  ?    State  some  results  produced  by  Moissan 
with  this  furnace.    Has  the  electric  furnace  any  industrial  use  ?    Where  ? 

10.  What  is  calcium  carbide?     How  is  it  made?     State  the  equation 
for  the  reaction.     What  are  its  properties ?     For  what  is  it  used? 

11.  What  is  carborundum?     How  is  it  made?     State  the  equation 
for  the  reaction.     What  are  its  properties  and  uses? 

12.  What  is  artificial  graphite?    How  is  it  made?    For  what  is  it  used? 

13.  Give  several  illustrations  of  the  production  of  (a)  electricity  by 
chemical  action,  and  (b)  vice  versa. 

14.  State  briefly  the  first  chemical  changes  which  were  produced  by 
electricity. 

15.  Describe  a  simple  voltaic  cell.     Why  is  it  so  called?    What  is 
the  source  of  the  electric  current  manifested  by  the  cell?     What  is  an 
electric  battery  ?     For  what  is  it  used  ? 

1 6.  Define  and  illustrate  (a)  electrolysis,  (b)  electrolyte,  (c}  elec- 
trode, (W)  anode,  (e)  cathode,  (/)  ions,  (g)  anion,  (h}  cation,  (*)  posi- 
tive electrode,  (/)  negative  electrode,  (£)  ionization. 

17.  Where  are  (a)  anions  and  (b)  cations  liberated? 

1 8.  Describe  an  electrolytic  cell.     How  does  it  differ  from  a  voltaic 
cell  ?     For  what  is  it  used  ? 

19.  Describe  the  electrolysis  of  (a)  zinc  chloride,  (b}  sodium  chloride, 
(c)  water. 

20.  State  and  illustrate  Faraday's  law. 

2 1 .  Give  a  brief  account  of  Faraday's  contribution  to  electrochemistry. 

22.  Describe  the  process  of  (a)  electrotyping  and  (b}  electroplating. 

23.  State  some  industrial  applications  of  the  electric  current. 

24.  What  is  the  theory  of  electrolysis?    What  is  the  present  theory 
of  solution  in  water?     What  is  the  theory  called?    Why?     What  facts 
support  it? 

25.  Define  and  illustrate  an  ionic  equation. 


Descriptive   Chemistry. 


PROBLEMS. 

1.  Calculate  the  percentage  composition  of  (a)  water,  (£)  magnetic 
oxide  of  iron  (Fe3O4),  (c}  crystallized   sodium  carbonate  (Na.,CO3  . 
ioH2O). 

2.  If  a  certain  current  of  electricity  deposited  31.7  gm.  of  copper, 
how  much   (a)   silver,  (&)   aluminium,  and   (c)    magnesium  would   it 
deposit  ? 

3.  If  a  certain  current  of  electricity  deposited  2  kg.  of  copper,  how 
much  silver  would  it  deposit? 

4.  How  much  calcium  carbide  can  be  made  (theoretically)  from  a 
ton  of  lime?     (Equation  is  3  C  +  CaO  =  CaC2  +  CO  or  36  +  56  = 
64  +  28.) 

5.  How  much  carborundum  can  be  made  (theoretically)  from  a  ton 
of  sand  (SiO2)  ?     (Equation  is  SiO2  +  3  C  =  SiC  +  2  CO  or  60  +  36 
=  40  +  56.) 

6.  Calculate  the  percentage  composition   of  (#)  carborundum  and 
(&)  calcium  carbide. 


CHAPTER    XL 
CHLORINE  AND  HYDROCHLORIC  ACID. 

CHLORINE  is  an  important  element,  and  its  compounds 
are  useful,  especially  hydrochloric  acid,  sodium  chloride, 
and  bleaching  powder. 

Occurrence.  —  Free  chlorine  is  never  found  in  nature, 
because  it  combines  so  readily  with  other  elements.  But 
in  combination  it  is  widely  distributed,  since  it  is  one  of 
the  components  of  common  salt,  or  sodium  chloride. 
Many  compounds  of  chlorine  with  potassium,  magnesium, 
and  calcium  are  found  in  the  deposits  at  Stassfurt  in  Ger- 
many (see  these  metals).  The  salts  found  in  sea  water 
contain  about  2  per  cent,  and  the  earth's  crust  contains 
about  o.oi  per  cent  of  chlorine.  Silver  chloride  —  "horn" 
silver  —  is  mined  as  an  ore  in  the  United  States  and 
Mexico. 

Preparation.  —  Chlorine  is  prepared  in  the  laboratory 
by  heating  a  mixture  of  manganese  dioxide  and  hydro- 
chloric acid.  This  method  was  used  by  Scheele,  who 
discovered  the  gas  in  1774.  The  equation  for  the  prepa- 
ration of  chlorine  is  — 

MnO2     +     4HC1     =     C12     +      MnCl2     +     2  H2O 
Manganese     Hydrochloric     Chlorine      Manganese  Water 

Dioxide  Acid  Bichloride 

V 

This  is  an  oxidizing  process,  since  the  hydrogen  of  the 
hydrochloric  acid  is  oxidized  to  water,  although  only  part 
of  the  chlorine  of  the  acid  is  obtained  free. 


134  Descriptive  Chemistry. 

Sometimes  chlorine  is  prepared  in  the  laboratory  by  heating  a  mixture 
of  manganese  dioxide,  sodium  chloride,  and  sulphuric  acid.  This  method 
is  substantially  the  same  as  the  other,  since  a  mixture  of  sulphuric  acid 
and  sodium  chloride  yields  hydrochloric  acid.  The  simplest  equation 
for  this  method  of  preparing  chlorine  is  — 

2  H2SO4  +  2  NaCl   +   MnO2    =     C12  +  Na,SO4  +   MnSO4  -f  2  H2O 
Sulphuric    Sodium  Manganese  Chlorine  Sodium    Manganese     Water 
Acid       Chloride     Dioxide  Sulphate     Sulphate 

Other  oxidizing  substances  besides  manganese  dioxide  may  be  used, 
such  as  potassium  chlorate  (KC1O3),  potassium  dichromate  (K2Cr2O7), 
and  red  lead  (Pb3O4). 

Chlorine  is  manufactured  by  several  processes,  all  of 
which  involve  the  same  principle  as  the  laboratory  method. 

In  the  Deacon  process,  hydrochloric  acid  is  oxidized  by  oxygen  ob- 
tained from  the  atmosphere.  A  mixture  of  hydrochloric  acid  gas  and 
air  is  heated  to  500°  C.  and  passed  through  iron  tubes  containing  balls 
of  clay  or  pieces  of  brick  previously  saturated  with  copper  chloride.  A 
series  of  complex  reactions  occurs  which  are  not  well  understood.  It  is 
supposed  that  the  copper  chloride  facilitates  the  formation  of  chlorine 
by  continuously  giving  and  taking  this  gas.  The  essential  chemical 
change,  however,  is  the  oxidation  of  the  hydrochloric  acid,  and  it  may 
be  represented  by  the  equation  — 

2HC1  -I-  O  C12  -f  H2O 

Hydrochloric  Acid      Oxygen  Chlorine  Water 

In  the  Weldon  process,  an  impure  native  manganese  dioxide,  known  as 
pyrolusite,  is  treated  with  hydrochloric  acid  in  large  earthenware  retorts 
or  stone  tanks  heated  by  hot  water  or  steam.  When  no  more  chlorine 
is  liberated,  the  residue  is  mainly  manganese  dichloride.  This  "  still- 
liquor"  was  formerly  thrown  away,  but  by  the  Weldon  process  it  is 
changed  into  manganese  compounds,  which  are  used  to  prepare  more 
chlorine  (see  Manganese  Dioxide). 

Chlorine  is  also  prepared  on  a  large  scale  by  the  electrolytic  process. 
Sodium  chloride  is  decomposed  by  electricity  in  properly  constructed 
cells,  and  the  chlorine  which  is  liberated  at  the  anode  is  conducted  off 
through  pipes  to  the  bleaching  powder  factory.  Sodium  hydroxide  is 
produced  at  the  same  time,  and  the  process  will  be  described  under  this 
compound. 


Chlorine  and  Hydrochloric  Acid.  135 

Properties.  —  Chlorine  is  a  greenish  yellow  gas.  Its 
color  suggested  the  name  chlorine  (from  the  Greek  word 
chloros,  meaning  greenish  yellow),  which  was  given  to  it 
by  Davy  about  1810.  It  has  a  disagreeable,  suffocating 
odor,  which  is  very  penetrating.  If  breathed,  it  irritates 
the  sensitive  lining  of  the  nose  and  throat,  and  a  large 
quantity  would  doubtless  cause  death.  It  is  heavier  than 
the  other  elementary  gases,  and  is  about  2.5  times  heavier 
than  air.  Hence  it  is  easily  collected  by  downward  dis- 
placement, i.e.  by  allowing  it  to  fall  to  the  bottom  of  a 
bottle  and  thus  fill  the  latter  by  displacing  the  air. 

A  liter  of  dry  chlorine  at  o°  C.  and  760  mm.  weighs  3.18  gm. 

Water  dissolves  chlorine.  The  solution  is  yellowish, 
smells  strongly  of  chlorine,  and  is  frequently  used  in  the 
laboratory  as  a  substitute  for  the  gas.  Chlorine  water,  as 
the  solution  is  called,  is  unstable  even  under  ordinary  con- 
ditions, and  must  be  kept  in  the  dark.  If  the  solution  is 
placed  in  the  sunlight,  oxygen  is  soon  liberated  and  hydro- 
chloric acid  is  formed.  Intermediate  changes  doubtless 
occur ;  but  the  simplest  equation  for  the  essential  change 

is-        Hp       +       C12       =       2HC1        +       O 

Water  Chlorine      Hydrochloric  Acid       Oxygen 

Chlorine  is  much  less  soluble  in  a  solution  of  sodium  chloride,  over 
which  it  is  sometimes  collected.  It  attacks  mercury  and  cannot  be  col- 
lected over  this  liquid. 

Chlorine  does  not  burn  in  the  air,  but  many  substances 
burn  in  chlorine.  The  metals  antimony  and  arsenic,  when 
sprinkled  into  chlorine,  suddenly  burst  into  flame,  while 
phosphorus  melts  at  first  and  finally  burns  with  a  feeble 
flame.  If  sodium,  iron  powder,  brass  wire,  or  other  metals 
are  heated  and  then  put  into  chlorine,  they  burn ;  the' 
sodium  and  iron  produce  a  dazzling  light  and  the  brass 


136  Descriptive  Chemistry. 

glows  and  emits  dense  fumes  of  whitish  smoke.  Chlorine 
combines  readily  with  hydrogen.  Hence,  a  jet  of  burning 
hydrogen  when  lowered  into  chlorine  continues  to  burn, 
forming  hydrochloric  acid  gas,  which  appears  as  a  white 
cloud.  The  simplest  equation  for  this  change  is  — 

H       +       Cl  HC1 

Hydrogen        Chlorine         Hydrochloric  Acid 

The  attraction  between  chlorine  and  hydrogen  is  so  great 
that  many  compounds  of  hydrogen  are  decomposed  by 
chlorine.  Thus,  compounds  containing  hydrogen  and 
carbon,  such  as  illuminating  gas,  paraffin  wax,  and  wood, 
burn  in  chlorine  with  a  smoky  flame.  Chlorine  does  not 
combine  directly  with  carbon,  hence  the  flame  consists 
largely  of  very  fine  particles  of  solid  carbon.  Similarly, 
a  piece  of  glowing  charcoal  is  extinguished  by  chlorine.  If 
filter  paper  is  saturated  with  warm  turpentine  (a  compound 
of  hydrogen  and  carbon)  and  put  into  a  bottle  of  chlorine, 
a  flame  accompanied  by  a  dense  cloud  of  black  smoke 
bursts  from  the  bottle  ;  the  chlorine  withdraws  the  hydro- 
gen to  form  hydrochloric  acid,  while  the  carbon  is  left  free. 
The  power  to  bleach  is  the  most  striking  and  useful 
property  of  chlorine.  This  property  depends  upon  the 
fact,  already  mentioned,  that  chlorine  withdraws  hydrogen 
and  liberates  free  oxygen  ;  the  latter  then  decomposes  the 
coloring  matter  in  the  cloth  or  other  material.  Dry 
chlorine  does  not  bleach.  If  an  envelope  on  which  the 
postmark,  or  a  lead  pencil  mark,  is  still  visible  is  placed 
in  moist  chlorine,  these  marks  will  not  be  bleached  be- 
cause they  are  largely  carbon ;  but  the  writing  ink,  which 
is  mainly  a  compound  of  hydrogen,  carbon,  and  iron,  will 
disappear.  Litmus  paper  and  calico  are  both  bleached  by 
moist  chlorine. 


Chlorine  and   Hydrochloric  Acid.  137 

Bleaching  Powder  is  the  source  of  the  chlorine  used  in 
the  bleaching  industries.  It  is  sometimes  called  "bleach," 
or  "  chloride  of  lime."  It  is  a  yellowish  white  substance 
having  a  peculiar  odor,  which  resembles  that  of  chlorine. 
When  dry,  it  is  a  powder,  but  on  exposure  to  the  air,  it 
absorbs  water  and  carbon  dioxide,  becomes  lumpy  and 
pasty,  and  loses  some  of  its  chlorine.  Acids  like  sulphuric 
and  hydrochloric  acid  liberate  from  bleaching  powder  its 
"  available  chlorine,"  which  varies  from  30  to  38  per  cent 
in  good  qualities.  The  equations  for  the  interaction  of 
acids  and  bleaching  powder  are  usually  written  thus  — 

CaOCl2         +       H2SO4    =    Cla    +    CaSO4    +    H2O 
Bleaching  Powder        Sulphuric  Acid  Calcium  Sulphate 

CaOCl2     +    2  HC1     =     C12     -f     CaCl2     +     H2O 
Hydrochloric  Acid  Calcium  Chloride 

The  composition  of  bleaching  powder  has  been  much  discussed. 
The  most  reliable  authority  gives  it  the  formula  CaOCL,.  When  dis- 
solved in  water,  bleaching  powder  forms  calcium  hypochlorite  (CaO2Cl2) 
and  calcium  chloride  (CaCl.,). 

Bleaching  Powder  is  manufactured  by  the  action  of  chlorine  gas  on 
lime.  Lime  (calcium  oxide,  CaO)  is  carefully  slaked  with  water  to 
form  calcium  hydroxide  (Ca(OH).2).  This  powder  is  sifted  into  a 
large  absorption  chamber  made  of  iron,  lead,  or  tarred  brick  until  the 
floor  is  covered  with  a  layer  three  or  four  inches  deep.  The  chlorine 
enters  at  the  top  and  settles  slowly  to  the  floor,  where  it  is  absorbed 
by  the  lime. 

The  simplest  equation  for  the  formation  of  bleaching  powder  might 
be  written  — 

Ca(OH)2          +          C12  CaOCl2        +        H2O 

Calcium  Hydroxide        Chlorine        Bleaching  Powder        Water 

Bleaching.  —  Immense  quantities  of  bleaching  powder 
are  used  to  whiten  cotton  and  linen  goods  and  paper  pulp. 
The  pieces  of  cotton  cloth  as  they  come  from  the  mill  are 


138  Descriptive  Chemistry. 

sewed  end  to  end  in  strips,  which  are  stamped  at  the 
extreme  ends  with  some  indelible  mark  to  distinguish  each 
owner's  cloth.  These  strips,  which  are  often  several  miles 
long,  are  drawn  by  machinery  into  and  out  of  numerous 
vats  of  liquors  and  water,  between  rollers,  and  through 
machines,  until  they  are  snow-white  and  ready  to  be 
finished  (i.e.  starched  and  ironed)  or  dyed.  The  whole 
operation  requires  three  or  four  days. 

The  preliminary  treatment  consists  in  singeing  off  the  downy  pile 
and  loose  threads  by  drawing  the  cloth  over  hot  copper  plates  or 
through  a  series  of  gas  flames.  The  object  of  the  remaining  operations 
is  threefold,  (i)  to  wash  out  mechanical  impurities,  the  fatty  and  resin- 
ous matter,  and  the  excess  of  the  different  chemicals,  (2)  to  remove 
matter  insoluble  in  water,  and  (3)  to  oxidize  the  coloring  matter  by 
chlorine.  The  details  of  the  process  differ  with  the  texture  of  the 
cloth  and  with  its  ultimate  use.  The  threefold  object  above  mentioned 
involves  successively  "liming,"  "souring,"  "chemicking,"  and  "souring,11 
interspersed  with  frequent  washing.  The  "liming11  consists  in  boiling 
the  cloth  in  a  large  kier  or  vat  with  lime,  the  "souring11  in  wetting  it 
with  weak  sulphuric  or  hydrochloric  acid,  and  the  "  chemicking ?1  in  im- 
pregnating it  with  a  weak  solution  of  bleaching  powder.  Often  the  cloth 
is  boiled  at  a  certain  stage  with  resin  and  sodium  carbonate.  The 
^liming11  removes  the  resinous  and  the  fatty  matter,  the  first  "souring11 
neutralizes  traces  of  lime,  and  the  second,  which  follows  the  "chem- 
icking,11 liberates  the  chlorine  in  the  fiber  of  the  cloth.  Frequent  washing 
is  absolutely  necessary  to  remove  the  impure  products  of  the  chemical 
changes  as  well  as  the  excess  of  lime  and  other  alkali,  acid,  and  chlo- 
rine. Should  these  be  left,  the  cloth  would  be  unevenly  bleached  and 
its  fiber  would  be  weak.  The  cloth  is  finally  treated  with  an  antichlor, 
such  as  sodium  hyposulphite,  which  removes  the  last  traces  of  chlorine. 

Bleaching  is  chemically  an  oxidizing  process.  The 
oxygen  when  it  is  liberated  from  water  by  chlorine  is  said 
to  be  in  the  nascent  state.  This  means  that  the  gas  is 
exceedingly  active,  because  it  is  not  only  uncombined,  but 
just  ready  to  unite  with  those  elements  for  which  it  has 
great  affinity.  Hence  this  nascent  oxygen  literally  tears 


Chlorine  and  Hydrochloric  Acid.  139 

down  complex  colored  substances  and  changes  them  into 
colorless  compounds.  The  nascent  state  is  aptly  illustrated 
by  bleaching  because  both  the  chlorine  and  the  oxygen 
are  in  this  active  chemical  condition. 

Chlorine  Hydrate  is  formed  by  cooling  chlorine  water  or  by  passing 
chlorine  into  ice  water.  It  is  a  yellowish,  crystalline  solid,  and  in  the 
air  it  decomposes  quickly  into  chlorine  and  water.  Its  composition 
corresponds  to  the  formula  C12  •  10  H2O. 

Liquid  Chlorine  was  first  prepared  by  Faraday  in  1823.  A  little 
chlorine  hydrate  was  inclosed  in  one  arm  of  a  bent  tube  (Fig.  21), 
which  was  then  sealed.  By  gently  heating  the  tube,  the  chlorine  hy- 
drate was  decomposed  into  chlorine  and  water, 
but  the  chlorine,  being  unable  to  escape,  was 
condensed  to  a  liquid  by  the  pressure  inside  the 
tube.  The  liquefaction  is  more  easily  accom- 
plished if  one  end  is  kept  cold  during  the 
experiment.  FlG-  21.-  Bent  tube  for 

.       .  ,.  ,,     .  .the   liquefaction  of   chlo- 

At  the  ordinary  pressure,  chlorine  gas  be-     rjne 

comes  liquefied,  if  its  temperature  is  —  34°  C, 

while  at  a  pressure  of  six  atmospheres  the  temperature  need  be  only 
o°  C.  Liquid  chlorine  has  a  bright  yellow  color.  It  is  a  commercial 
article,  and  is  stored  and  shipped  in  steel  cylinders  lined  with  lead. 
It  is  used  in  the  laboratory  to  prepare  chlorides,  and  industrially  to 
extract  gold.  Solid  chlorine  has  been  obtained  as  a  yellow  crystalline 
mass  by  cooling  the  liquid  to  —  102°  C. 

Uses  of  Chlorine.  —  Chlorine  is  used  directly  to  prepare 
some  of  its  compounds,  the  most  important  being  bleaching 
powder.  The  latter  is  often  used  as  a  deodorizer  and  dis- 
infectant, since  the  liberated  chlorine  destroys  putrefying 
matter  by  acting  on  it  as  on  coloring  matter.  A  solution 
of  potassium  hypochlorite  (Javelle's  water)  or  sodium  hy- 
pochlorite  (Labarraque's  solution)  is  often  used  to  remove 
fruit  stains  from  cotton  and  linen  goods. 

Chlorides  are  formed  when  chlorine  combines  with  other 
elements,  and  they  are  in  general  stable  compounds. 


140  Descriptive  Chemistry. 

The  simplest  equations  illustrating  the  combination  of  chlorine  with 
metals  and  other  elements  are  — 

Na  +       Cl       =          NaCl 

Sodium  Chlorine      Sodium  Chloride 

Sb          +     3d      =  SbCl3 

Antimony  Antimony  Trichloride 

Cu  +       C12      =  CuCl2 

Copper  Copper  Chloride 

P  +     3C1      =  PC13 

Phosphorus  Phosphorus  Trichloride 

H  +       Cl       =  HC1 

Hydrogen  Hydrochloric  Acid 

Chlorides  form  an  important  class  of  compounds  and  they  will  be 
considered  under  the  elements  with  which  the  chlorine  combines. 
(See  also  Chlorides  below.) 

HYDROCHLORIC    ACID. 

Hydrochloric  Acid  is  the  most  useful  compound  of 
chlorine.  It  is  a  gas,  very  soluble  in  water^  This  solution 
has  long  been  known  as  muriatic  acid  (from  the  Latin 
word  muria,  meaning  brine).  The  term  hydrochloric  acid 
includes  both  the  gas  and  its  solution,  but  the  solution  is 
usually  meant. 

The  early  chemists  called  the  gas  "  spirit  of  salt."  Priestley,  who 
first  prepared,  collected,  and  studied  the  gas,  called  it  "  marine  acid  air." 
Both  expressions  emphasize  its  relation  to  salt  (sodium  chloride). 

Occurrence.  —  The  gas  occurs  free  in  volcanic  gases. 
The  solution  is  one  constituent  of  the  gastric  juice  of  the 
stomach.  Chlorides,  which  are  salts  of  hydrochloric  acid, 
are  abundant  in  the  earth's  crust. 

Preparation. — The  gas  is  prepared  in  the  laboratory 
by  the  method  devised  by  Glauber  in  the  seventeenth  cen- 


Chlorine  and  Hydrochloric  Acid.  141 

tury,  viz.,  by  heating  sulphuric  acid  and  sodium  chloride. 
If  the  mixture  is  gently  heated,  the  chemical  change  is 
represented  thus  — 

Nad  +  H2S04  =  HC1  +  HNaSO4 
Sodium  Sulphuric  Hydrochloric  Acid  Sodium 
Chloride  Acid  Acid  Sulphate 

But  at  a  high  temperature  the  equation  for  the  reaction 
2  NaCl  +  H2SO4  -  2  HC1  +  Na2SO4 

In  either  case  the  gas  is  readily  produced.  It  may  be 
collected  over  mercury  or,  more  easily,  by  downward  dis- 
placement. The  solution  is  prepared  by  passing  the  gas 
into  water. 

That  sodium  sulphate  is  the  other  product  of  the  chemical  change  at 
a  high  temperature  may  be  shown  by  testing  the  heated  residue  as 
follows  :  (a)  Dissolve  a  portion  in  water  and  add  a  few  drops  of  barium 
chloride  solution ;  the  immediate  formation  of  the  white,  insoluble 
barium  sulphate  shows  that  the  residue  from  the  experiment  must  be 
a  sulphate,  (b}  Burn  a  little  of  the  residue  on  a  platinum  wire  or 
piece  of  porcelain  held  in  the  Bunsen  flame ;  the  intense  yellow  color 
immediately  imparted  to  the  flame  shows  that  the  residue  contains 
sodium,  (c)  Hence  the  compound  must  be  sodium  sulphate. 

Commercial  Hydrochloric  Acid  is  manufactured  in  enor- 
mous quantities  by  the  method  used  in  the  laboratory. 
A  mixture  of  salt  and  sulphuric  acid  is  moderately  heated 
in  a  large  hemispherical  cast-iron  pan,  and  the  gas  passes 
through  an  earthenware  pipe  into  an  absorbing  tower ;  the 
fused  mass  of  acid  sodium  sulphate  and  salt  is  then  sub- 
jected to  a  higher  temperature,  and  the  liberated  gas  passes 
by  another  pipe  into  the  absorbing  tower.  These  towers 
are  tall  and  filled  with  coke  or  pieces  of  brick  over  which 
water  trickles  ;  as  the  hydrochloric  acid  gas  passes  up  the 
tower,  it  is  absorbed  by  the  descending  water,  and  flows 


142  Descriptive  Chemistry. 

out  at  the  bottom  of  the  tower  as  concentrated  acid.  The 
gas  is  usually  cooled  before  it  enters  the  towers.  Some- 
times the  gas  passes  through  huge  earthenware  jars  be- 
fore entering  the  towers.  In  these  jars  the  gas  and  water 
are  caused  to  flow  constantly  in  opposite  directions,  thus 
insuring  complete  absorption. 

Hydrochloric  acid  gas  is  a  by-product  in  the  manufacture  of  sodium 
carbonate  by  the  Leblanc  process.  The  gas  was  formerly  allowed  to 
escape  into  the  atmosphere,  but  since  it  destroyed  vegetation  and  be- 
came a  nuisance  in  other  ways,  a  law  was  passed  forbidding  the  manu- 
facturers to  let  it  escape.  Hence  it  became  necessary  to  absorb  the 
gas  in  water.  The  hydrochloric  acid,  which  was  once  regarded  as  a 
waste  product,  is  now  the  main  source  of  profit,  since  competition  has 
reduced  the  price  of  sodium  carbonate  (see  Sodium  Carbonate). 

Properties.  —  Hydrochloric  acid  gas  is  colorless  and 
transparent.  When  it  escapes  into  moist  air,  it  forms 
fumes  which  are  really  minute  drops  of  a  solution  of  the 
gas  in  the  moisture  of  the  air.  It  has  a  choking,  sharp, 
pungent  odor.  The  gas  does  not  burn  nor  support  com- 
bustion. It  is  about  1.25  times  heavier  than  air,  and  may 
therefore  be  collected  by  downward  displacement. 

One  liter  at  o°C.  and  760  mm.  weighs  1.61  gm.  The  gas  can  be 
liquefied  at  io°C.  and  40  atmospheres  pressure;  while  at  —  i6°C,  the 
pressure  need  be  only  20  atmospheres. 

The  extreme  solubility  of  hydrochloric  acid  gas  in  water  is 
one  of  its  most  striking  properties.  One  liter  of  water  will 
dissolve  about  500  1.  of  gas,  if  both  are  at  o°  C.  and  760  mm. 
At  the  ordinary  temperature  about  450 1.  of  gas  dissolve  in 
i  1.  of  water,  and  as  the  temperature  rises  the  solubility 
decreases.  The  solution  is  the  familiar  hydrochloric  acid. 
The  gas  readily  escapes,  hence  the  acid  forms  fumes  when 
exposed  to  air.  Pure  hydrochloric  acid  is  a  colorless  liquid. 
The  commercial  acid  has  a  yellow  color,  usually  due  to  iron 


Chlorine  and  Hydrochloric  Acid.  143 

compounds,  but  sometimes  to  organic  matter  or  to  dissolved 
chlorine.  It  also  contains  other  impurities.  Like  most 
acids,  it  reddens  blue  litmus,  and  gives  up  its  hydrogen 
when  added  to  metals. 

The  strongest  acid  contains  about  42  per  cent  (by  weight)  of  the 
gas,  and  its  specific  gravity  is  i  .2.  When  the  strong  acid  is  heated,  the 
gas  is  evolved  until  the  solution  contains  about  20  per  cent  of  the  acid, 
and  then  the  liquid  boils  at  i  io°C.  without  further  change.  The  dilute 
acid,  on  the  other  hand,  loses  water  until  the  same  conditions  prevail. 

Composition  of  Hydrochloric  Acid  Gas.  —  In  1810,  Davy  showed 
that  hydrochloric  acid  gas  (which  had  been  regarded  as  an  oxygen 
compound)  contained  only  chlorine  and  hydrogen.  Many  facts  lead 
us  to  conclude  that  hydrochloric  acid  gas  is  composed  of  hydrogen  and 
chlorine  in  such  a  ratio  that  its  composition  is  represented  by  the  for- 
mula HC1.  (i)  Hydrogen  burns  in  chlorine,  and  the  only  product  is 
hydrochloric  acid  gas.  (2)  When  hydrochloric  acid  is  decomposed 
by  an  electric  current,  equal  volumes  of  hydrogen  and  chlorine  are 
evolved.  (3)  When  a  mixture  of  equal  volumes  of  hydrogen  and 
chlorine  is  exposed  to  the  direct  sunlight  or  to  the  action  of  an  electric 
spark,  the  gases  combine  with  an  explosion,  and  hydrochloric  acid  gas 
is  formed  with  no  residue.  Furthermore,  the  volume  of  the  resulting 
gas  equals  the  sum  of  the  volumes  of  hydrogen  and  chlorine  used. 
(4)  When  a  given  volume  of  dry  hydrochloric  acid  gas  is  treated  with 
sodium  amalgam,  the  chlorine  is  withdrawn  by  the  sodium  in  the  amal- 
gam, and  a  volume  of  hydrogen  remains  which  is  half  the  original  vol- 
ume. (5)  No  derivative  of  hydrochloric  acid  is  known  which  contains 
less  hydrogen, or  less  chlorine  in  a  molecule.  (6)  The  ratio  by  weight 
in  which  hydrogen  and  chlorine  combine  is  1:35.45.  Hence,  the 
lowest  molecular  weight  of  hydrochloric  acid  is  36.45,  a  number  which 
has  been  verified  by  several  different  methods. 

Uses  of  Hydrochloric  Acid.  —  Vast  quantities  are  used 
to  prepare  chlorine  for  the  manufacture  of  bleaching  pow- 
der. Various  chlorides  are  prepared  from  it,  and  it  is  one 
of  the  common  acids  used  in  chemical  laboratories. 

Chlorides  are  formed  by  the  direct  addition  of  chlorine 
to  metals,  as  we  have  seen.  They  are  also  formed  when 


144  Descriptive  Chemistry. 

metals,  their  oxides,  or  hydroxides  are  added  to  hydro- 
chloric acid.  The  following  equations  illustrate  this  gen- 
eral fact :  — 

Zn  +   2  HC1         =   ZnCl2  +     H2 

Zinc  Zinc  Chloride 

ZnO         +   2  HC1         =   ZnCl2  +     H2O 

Zinc  Oxide  Zinc  Chloride 

Zn(OH)2      +   2  HC1         -   ZnCl2  +    2  H2O 

Zinc  Hydroxide  Zinc  Chloride 

They  are  also  formed  by  adding  other  salts  to  hydro- 
chloric acid. 

Molecules  of  chlorides  may  contain  several  atoms  of  chlorine. 
Occasionally  the- name  of  the  compound  indicates  this  fact,  e.g.  manga- 
nese dichloride  (MnQ2),  antimony  trichloride  (SbCl;5),  phosphorus 
trichloride  and  pentachloride  (PC13  and  PCI-)-  If  a  metal  forms  two 
chlorides,  the  two  are  distinguished  .by  modifying  the  name  of  the 
metal.  The  one  containing  the  smaller  proportion  of  chlorine  ends  in 
-ous,  the  one  containing  the  larger  ends  in  -ic.  Thus,  mercurous  chlo- 
ride is  HgCl,  but  HgCl2  is  mercuric  chloride.  Similarly,  we  have  fer- 
rous chloride,  FeCl2,  and  ferric  chloride,  FeCl3. 

The  Test  for  Hydrochloric  Acid  and  Chlorides.  — Most 

chlorides  are  soluble  in  water.  Those  of  lead,  silver,  and 
mercury  (-ous)  are  not.  If  silver  nitrate  is  added  to  hydro- 
chloric acid,  or  to  the  solution  of  a  chloride,  a  white,  curdy 
precipitate  of  silver  chloride  is  formed,  which  (a)  is  insol- 
uble in  nitric  acid,  but  soluble  in  warm  ammonium  hydrox- 
ide, and  (£)  turns  purple  in  the  sunlight.  The  invariable 
formation  of  silver  chloride  is  the  test  for  hydrochloric 
acid  and  soluble  chlorides.  Hydrochloric  acid  gas  also 
forms  dense  white  clouds  of  ammonium  chloride  in  the 
presence  of  ammonia  gas. 


Chlorine  and    Hydrochloric  Acid.  145 

Miscellaneous.  —  The  acids  of  chlorine  are  tabulated  under  ACIDS. 
The  compounds  of  chlorine  with  sodium,  potassium,  magnesium,  and 
calcium  are  described  under  these  metals. 

Aqua  regia,  of  which  chlorine  is  one  constituent,  is  discussed  in 
Chapter  XII. 

EXERCISES. 

1.  What  is  the  symbol  of  chlorine  ?     What  useful  compounds  con- 
tain this  element  ? 

2.  How  is  chlorine  prepared  in  the  laboratory  ?     Give  one  equation 
for   its    preparation.      Describe    Deacon^   process   for    manufacturing 
chlorine. 

3.  Who  discovered  chlorine  ?     Who  named  it,  when,  and  why  ? 

4.  Summarize  the  physical  properties  of  chlorine.      How  can  it  be 
quickly  distinguished  from  the  gases  previously  studied  ? 

5.  Summarize   the   chemical    properties  of  chlorine.      Compare  it 
with  oxygen.     Describe  fully  its  action  with  hydrogen. 

6.  Define    (a)    downward    displacement,     (b}    available    chlorine, 
(V)  antichlor. 

7.  Develop  the  topics :  (a)  nascent  state,  (<$)  chlorine  water,  (V)  chlo- 
rine hydrate,  (ti )  liquid  chlorine,  (e)  chlorine  is  an  oxidizing  agent. 

8.  What  is  bleaching  powder  ?     How  is  it  made  ?     What  are  its 
chief  properties  ?     Describe  the  operation  of  bleaching.     What  is  the 
chemistry  of  bleaching  ? 

9.  What  is  (a)  "bleach,"  (£)  muriatic  acid,  (V)  chloride  of  lime, 
(d)  "salt,11  (e)  "lime,"  (/)  commercial  hydrochloric  acid  ? 

10.  What  are  chlorides  ?     Name  five.     How  can  they  be  formed  ? 
Give  the  formula  of  sodium  chloride.    Why  cannot  chlorine  be  collected 
over  mercury  ? 

11.  What  is  hydrochloric  acid  ?     How  is  it  prepared  in  the  labora- 
tory ?     Give   the   equations  for  its  preparation.     How  is  it  prepared 
industrially  ? 

12.  Summarize  the  chief  properties  of  hydrochloric  acid  gas.    Of  the 
acid,  as  the  term  is  usually  used.     What  happens  when  hydrochloric 
acid  is  boiled  ? 

13.  What  is  the  evidence  that  the  formula  of  hydrochloric  acid  gas 
is  HC1  ? 

14.  For  what  is  hydrochloric  acid  used  ?     State  the  test  for  hydro- 
chloric acid  and  soluble  chlorides. 


146  Descriptive  Chemistry. 

15.  Give  a  brief  account  of  Faraday's  work  on  chlorine.     Of  Davy's 
work. 

1 6.  Why  is  chlorine  never  found  free  ? 


PROBLEMS. 

1.  One  equation  for  the  preparation  of  chlorine  is  — 

4HC1     +     MnO2      =   C12      +     MnCI2      +     2H2O 
146       +        87         =71       +       126         +         36 

(0)    How  many  grams  of  chlorine  can  be  made  from  247  gm.  of  man- 
ganese dioxide  ?     (£)   Name  all  the  products. 

2.  How  much  sodium  chloride  is  needed  to  prepare  a  kilogram  of 
hydrochloric  acid  gas  ? 

3.  How  many  grams  of  manganese  dioxide  are  necessary  to  ^prepare 
100  gm.  of  chlorine  from  hydrochloric  acid. 

4.  A  bottle  of  chlorine  water  was  exposed  to  the  sunlight  until 
all  the  chlorine  disappeared,     (a)  What  two  products  were  formed  ? 
(^)  Write  the  equation  for  the  reaction.     (c}    What  weight  of  chlorine 
gas  is  necessary  to  form  20  gm.  of  the  gaseous  product  ?      (d)  What 
volume  of  chlorine  is  necessary  to  form  20  gm.  of  the  other  product  ? 

5.  Calculate  the  percentage  composition  of  (a)  hydrochloric  acid 
gas,  (b)  sodium  chloride,  (c)  silver  chloride  (AgCl),  (d)  potassium 
chloride  (KC1). 


CHAPTER   XII. 
COMPOUNDS  OF  NITROGEN. 

THE  most  important  compounds  of  nitrogen  are  am- 
monia (NH3),  nitric  acid  (HNO3),  and  compounds  related 
to  them.  Many  animal  and  vegetable  substances  essential 
to  life  are  compounds  of  nitrogen. 

AMMONIA. 

The  term  ammonia  includes  both  the  gas  and  its  solu- 
tion in  water,  though  the  latter  is  more  accurately  called 
ammonium  hydroxide. 

Formation  of  Ammonia.  —  When  vegetable  and  animal 
matter  containing  nitrogen  decays,  the  nitrogen  and  hydro- 
gen are  liberated  in  combination,  as  ammonia.  The  odor 
of  ammonia  is  often  noticed  near  stables.  If  animal  sub- 
stances containing  nitrogen  are  heated,  ammonia  is  given 
off.  The  old  custom  of  preparing  ammonia  by  heating 
horns  and  hoofs  in  a  closed  vessel,  i.e.  by  dry  distillation, 
gave  rise  to  the  term  "spirits  of  hartshorn."  Soft  coal 
contains  compounds  of  nitrogen  and  of  hydrogen,  and  when 
the  coal  is  heated  to  make  illuminating  gas,  one  of  the  prod- 
ucts is  ammonia. 

Preparation. — Ammonia  gas  is  prepared  in  the  labora- 
tory by  heating  ammonium  chloride  with  an  alkali,  usually 
slaked  lime.  The  reaction  may  be  represented  thus  — 

H7 


148  Descriptive  Chemistry. 

2NH4C1   +  Ca(OH)2  =    2NH3     +    CaCl2  +  2  H2O 
Ammonium  Slaked  Ammonia  Calcium 

Chloride  Lime  Gas  Chloride 

107       +       74         =34         +       111    +    36 
The  gas  is  usually  collected  by  upward  displacement,  i.e. 
by  allowing  the  gas  to  flow  upward  into  a  bottle  and  dis- 
place the  air.     The  solution  is  prepared  by  conducting  the 
gas  into  water. 

The  main  source  of  the  ammonia  of  commerce  is  the  ammoniacal 
liquor  or  gas  liquor  of  the  gas  works.  The  gases  which  come  from  the 
retorts  in  which  the  coal  is  heated  are  passed  into  water,  which  absorbs 
the  ammonia  and  some  other  gases.  This  impure  gas  liquor  is  treated 
with  lime  to  liberate  the  ammonia,  which  is  absorbed  in  tanks  contain- 
ing hydrochloric  acid  or  sulphuric  acid.  This  solution  upon  the  addi- 
tion of  an  alkali  gives  up  its  ammonia,  which  is  dissolved  in  distilled 
water,  forming  thereby  the  ammonium  hydroxide  or  aqua  ammonia 
of  commerce. 

Ammonia  is  sometimes  prepared  from  the  residues  of  the  beet  sugar 
industry,  from  the  refuse  of  slaughter  houses  and  tanneries,  and  from  the 
gases  from  coke  ovens.  It  is  not  obtained  directly  from  the  nitrogen  of 
the  air. 

Properties  of  Ammonia.  —  Ammonia  gas  is  colorless. 
It  has  an  exceedingly  pungent  odor,  and  if  inhaled  sud- 
denly or  in  large  quantities  it  brings  tears  to  the  eyes  and 
may  cause  suffocation.  It  is  a  light,  volatile  gas,  being  only 
.59  times  as  heavy  as  air.  A  liter  of  the  gas  at  o°  and 
760  mm.  weighs  .77  gm.  It  will  not  burn  in  the  air,  nor 
will  it  support  the  combustion  of  a  blazing  stick  ;  but  if  the 
air  is  heated  or  if  its  proportion  of  oxygen  is  increased,  a 
jet  of  ammonia  gas  will  burn  in  it  with  a  yellowish  flame, 
thereby  illustrating  the.  broader  application  of  the  term 
combustion. 

Ammonia  gas  is  easily  liquefied  if  reduced  to  o°C.  and 
subjected  to  a  pressure  of  4^  atmospheres,  while  at  —  34°  C. 
it  liquefies  at  the  ordinary  atmospheric  pressure. 


Compounds  of  Nitrogen.  149 

Liquefied  ammonia  is  often  called  anhydrous  ammonia,  because  it 
contains  no  water.  It  boils  at  —  33. 5° C.  Hence,  if  it  is  exposed  to 
the  air  or  warmed  in  any  way,  it  changes  back  to  a  gas,  and  in  so  doing 
absorbs  considerable  heat.  This  fact  has  led  to  the  extensive  use  of 
liquid  ammonia  in  the  manufacture  of  ice. 

Ammonia  is  a  strong  alkali,  and  was  called  formerly  the 
volatile  alkali.  Priestley,  who  discovered  and  studied  the 
gas,  called  it  alkaline  air. 

Another  marked  property  of  ammonia  gas  is  its  solu- 
bility in  water.  A  liter  of  water  at  o°C.  dissolves  1148!. 
of  gas  (measured  at  O°C.  and  760  mm.),  and  at  the 
ordinary  temperature  I  1.  of  water  dissolves  about  700 1.  of 
gas.  This  solution  of  the  gas  is  usually  called  ammonia, 
though  other  names,  especially  ammonium  hydroxide,  are 
sometimes  applied  to  it.  Commercially  it  is  known  as 
aqua  ammonia,  ammonia,  or  ammonia  water.  It  gives  off 
the  gas  freely,  when  heated,  as  may  easily  be  discovered 
by  the  odor  or  by  the  formation  of  the  dense  white  fumes 
of  ammonium  chloride  (NH4C1)  when  the  solution  is  ex- 
posed to  hydrochloric  acid.  The  solution  is  lighter  than 
water,  its  specific  gravity  being  about  .88,  and  contains 
about  35  per  cent  (by  weight)  of  the  gas.  It  is  a  strong 
alkali  —  a  caustic  alkali,  neutralizes  acids  and  forms  salts, 
and  acts  in  many  respects  like  sodium  hydroxide. 

Ammonium  Hydroxide  and  Ammonium  Compounds.— 

When  ammonia  gas  is  passed  into  water^it  is  believed  that 
the  ammonia  combines  with  the  water  and  forms  a  solution 
of  an  unstable  compound  having  the  formula  NH4OH. 
This  compound  is  ammonium  hydroxide  (or  ammonium 
hydrate).  Its  formation  may  be  represented  thus  — 

NH3         -f-     H2O  NH4OH 

Ammonia  Water  Ammonium  Hydroxide 


150  Descriptive  Chemistry. 

Ammonium  -hydroxide  acts  like  a  base.  It  has  a  marked 
alkaline  reaction ;  it  neutralizes  acids  and  forms  salts, 
thus — 

NH4OH      +       HC1  NH4C1         +      H2O 

Ammonium  Chloride 

2NH4OH     +     H2S04      =   (NH4)2S04      +     2  H2O 

Ammonium  Sulphate 

These  salts,  ammonium  chloride  and  ammonium  sul- 
phate, have  definite  properties,  and  are  strictly  analogous 
to  sodium  salts.  Thus,  we  have  — 

Sodium  Salts  Ammonium  Salts 
Nad  NH4C1 

NaNO3 *  NH4NO3 

Na2S04  (NH4)2S04 

etc.  etc. 

Hence,  it  is  believed  that  ammonium  compounds  contain  a 
group  of  atoms  which  acts  like  an  atom  of  a  metal.  This 
group  of  atoms  is  called  ammonium,  and  its  formula  is 
NH4.  Ammonium  has  never  been  separated  from  its 
compounds,  or  if  it  has  it  is  so  unstable  that  it  immedi- 
ately decomposes  into  ammonia  gas  and  hydrogen.  So 
also  ammonium  hydroxide  has  never  been  obtained  free, 
for  it  decomposes  readily  into  ammonia  gas  and  water, 

thus  — 

NH4OH  NH3      +  H2O 

Ammonium  Hydroxide      Ammonia  Gas      Water 

Ammonium  is  sometimes  called  a  radical,  because  it  is  the  root  or 
foundation  of  a  series  of  compounds.  It  is  likewise  called  a  hypotheti- 
cal metal,  because  its  existence  is  assumed  and  it  acts  chemically  like 
metals. 


Compounds  of  Nitrogen.  151 

Ammonium  Chloride  is  prepared  by  passing  ammonia 
gas  into  dilute  hydrochloric  acid,  by  mixing  ammonium 
hydroxide  and  hydrochloric  acid,  or  by  letting  the  two 
gases  mingle.  The  equation  for  the  essential  reaction  is — 

NH3    +  HC1  =  NH4C1 

Ammonia      Hydrochloric  Acid      Ammonium  Chloride 

It  is  convenient  to  regard  this  compound  as  the  ammonium 
salt  of  hydrochloric  acid,  as  if  it  were  formed  by  replacing 
the  hydrogen  of  the  acid  by  ammonium,  just  as  sodium 
forms  sodium  chloride. 

Ammonium  chloride  is  a  white,  granular  or  crystalline 
solid,  with  a  sharp,  salty  taste.  It  dissolves  easily  in 
water,  and  in  so  doing  lowers  the  temperature  markedly. 
When  heated  to  a  high  temperature  it  gradually  breaks 
up  into  ammonia  and  hydrochloric  acid.  This  kind  of 
decomposition  is  called  dissociation. 

Large  quantities  of  ammonium  chloride  are  made  at  one  stage  of  the 
manufacture  of  ammonium  hydroxide  by  passing  the  gas  into  hydro- 
chloric acid.  The  crude  product  is  called  "  muriate  of  ammonia  "  to 
indicate  its  relation  to  muriatic  (or  hydrochloric)  acid.  It  is  largely 
used  for  charging  Leclanche'  batteries,  as  an  ingredient  of  soldering 
fluids,  in  galvanizing  iron,  and  in  textile  industries.  The  crude  salt  is 
purified  by  heating  it  gently  in  a  large  iron  or  earthenware  pot,  with  a 
dome-shaped  cover ;  the  ammonium  chloride  volatilizes  easily  and  then 
crystallizes  in  the  pure  state  as  a  fibrous  mass  on  the  inside  of  the  cover, 
but  the  impurities  remain  behind  in  the  vessel.  The  process  of  vapor- 
izing a  solid  substance  and  then  condensing  the  vapor  directly  into 
the  solid  state  is  called  sublimation.  It  differs  from  distillation  in  that 
the  substance  does  not  pass  through  an  intermediate  liquid  state.  The 
product  of  sublimation  is  called  a  sublimate.  Sublimed  ammonium 
chloride  is  known  as  sal  ammoniac. 

Ammonium  Sulphate  is  made  by  passing  ammonia  gas  into  sul- 
phuric acid,  or  by  adding  ammonium  hydroxide  to  the  acid,  thus  — 

2NH4OH       +       H2SO4       =       (NH4)2SO4       +       2  H2O 
Ammonium  Hydroxide  Ammonium  Sulphate 


152  Descriptive  Chemistry. 

«*  The  commercial  salt  is  a  grayish  or  yellowish  solid.     It  is  used  as  a 
Constituent  of  fertilizers,  since  it  is  rich  in  nitrogen,  and  in  making 
ammonium  alum  and  other  ammonium  compounds. 

Ammonium  Nitrate  is  made  by  passing  ammonia  into  nitric  acid,  or 
by  allowing  ammonia  gas  and  the  vapor  of  nitric  acid  to  mingle,  thus  — 

NH3      +      HNO3      =  NH4NO3 

Ammonia     Nitric  Acid     Ammonium  Nitrate 

It  is  a  white  salt  which  forms  beautiful  crystals.  It  dissolves  easily  in 
water  with. a  fall  of  temperature.  Its  chief  use  is  in  the  preparation  of 
nitrous  oxide  (see  this  compound). 

Ammonium  Carbonate  is  an  impure  salt  as  found  in  commerce, 
being  a  mixture  of  acid  ammonium  carbonate  (HNH4CO3)  and  a 
related  compound.  When  pure  and  fresh  it  is  transparent,  but  on  ex- 
posure to  the  air  it  loses  ammonia  and  turns  white.  It  is  used  to  pre- 
pare some  kinds  of  baking  powder,  to  scour  wool,  as  a  medicine,  and 
to  prepare  smelling  salts,  since  it  gives  off  ammonia  readily. 

Other  ammonium  compounds  are  sodium  ammonium  phosphate 
or  microcosmic  salt  (HNaNH4PO4),  ammonium  sulphocyanate 
(NH4SCN),  and  ammonium  sulphide  ( (NH4)2S). 

Uses  of  Ammonia. — Ammonia  in  the  different  forms  is 
widely  used  as  a  cleansing  agent,  especially  for  the  re- 
moval of  grease,  as  a  restorative  in  cases  of  fainting  or  of 
inhaling  irritating  gases,  in  dyeing  and  calico  printing,  and 
in  the  manufacture  of  dyestuffs,  sodium  carbonate,  and 
ice.  Its  salts  have  many  domestic,  industrial,  and  agri- 
cultural uses. 

The  Use  of  Ammonia  as  a  Refrigerant  and  in  making 
Ice  depends  upon  the  fact  that  many  liquids  in  passing 
into  a  gas  absorb  heat.  Liquefied  ammonia  (not  the  ordi- 
nary liquid  ammonia)  changes  rapidly  into  a  gas  when  its 
temperature  is  raised  or  the  pressure  reduced.  Hence,  if 
anhydrous  ammonia  is  allowed  to  flow  through  a  pipe  sur- 
rounded by  brine,  the  ammonia  evaporates  in  the  pipe  and 
cools  the  brine,  which  may  be  used  as  a  refrigerant  or  for 


Compounds  of  Nitrogen.  153 

making  ice.  In  some  cold  storage  houses,  breweries, 
packing  houses,  and  sugar  refineries,  this  cold  brine  is 
pumped  through  pipes  placed  in  the  rooms  where  a  low 
temperature  is  desired. 

The  construction  and  operation  of  an  ice-making  plant  are  essentially 
as  follows  :  — 

Liquefied  ammonia  is  forced  from  a  tank  into  a  series  of  pipes  which 
are  submerged  in  an  immense  vat  filled  with  brine.  Large  galvanized 
iron  cans  containing  pure  water  to  be  frozen  are  immersed  in  the  brine, 
which  is  being  kept  below  the  freezing  point  of  water  by  the  rapid  evap- 
oration of  the  ammonia  in  the  pipes.  In  about  sixty  hours  the  water 
in  the  cans  is  changed  into  a  cake  of  ice  weighing  about  three  hundred 
pounds.  As  fast  as  the  ammonia  gas  forms  in  the  pipes,  it  is  removed 
by  exhaust  pumps  into  another  tank,  where  it  is  recondensed  to  liquefied 
ammonia  and  conducted,  as  needed,  into  the  first  tank  to  be  used  again. 
The  ammonia  is  thus  used  over  and  over  without  appreciable  loss. 
The  pure  water  is  sometimes  obtained  by  condensing  the  exhaust 
steam  from  the  boilers  used  to  operate  the  machinery,  though  it 
usually  comes  from  a  deep  well.  Most  ocean  steamers  have  an  ice 
plant,  and  in  large  cities  in  warm  climates  manufactured  ice  is  a  com- 
mon commodity. 

Composition  of  Ammonia  Gas.  — Numerous  experiments  show  that 
ammonia  gas  has  the  composition  expressed  by  the  formula  NH3. 

(1)  Dry  ammonia  gas  passed  over  heated  magnesium  decomposes  into 
hydrogen  and  nitrogen.     The  hydrogen  may  be  collected  and  tested, 
but  the  nitrogen  combines  with  the  magnesium,  forming  a  yellowish 
green  powder  called  magnesium  nitride,  thus  — 

2NH3     +     3Mg       =  Mg,N2  +     3H2 

Magnesium      Magnesium  Nitride 

These   facts   show   that    ammonia    contains    nitrogen   and    hydrogen. 

(2)  If  a  bottle  is  filled  with  chlorine  gas  and  plunged  mouth  downward 
into  a  vessel  containing  ammonium  hydroxide,  dense  white  fumes  fill 
the  bottle,  the  greenish  chlorine  gas  disappears,  and  the  liquid  rises  in 
the  bottle  ;  after  the  bottle  has  stood  mouth  downward  in  a  dish  con- 
taining dilute  hydrochloric  acid  (to  neutralize  the  excess  of  ammonia), 
the  gas  in  the  bottle  will  be  found  to  be  nitrogen.     The  chlorine  with- 


154  Descriptive  Chemistry. 

draws  the  hydrogen  from  the  ammonia  of  the  ammonium  hydroxide, 
leaving  the  nitrogen  free,  thus  — 

NH,      +      3  Q      -        N        +  3HC1 

Ammonia      Chlorine      Nitrogen      Hydrochloric  Acid 

(3)  The  same  experiment,  if  performed  accurately,  shows  that  one 
volume  of  nitrogen  combines  with  three  volumes  of  hydrogen  to  form 
ammonia  gas.  A  tube  containing  a  known  volume  of  chlorine  is  pro- 
vided with  a  funnel  through  which  concentrated  ammonium  hydroxide 
is  dropped  into  the  chlorine,  until  the  reaction  ceases  (Fig.  22).  After 
the  excess  of  ammonia  is  neutralized  with  sulphuric  acid,  the  volume 
of  nitrogen  left  is  one  third  of  the  original  volume 
of  chlorine  gas.  Now  hydrogen  and  chlorine  com- 
bine in  equal  volumes,  hence  the  volume  of  hydrogen 
withdrawn  from  the  added  ammonia  must  be  equal 
to  the  original  volume  of  chlorine.  But  this  volume 
is  three  times  the  volume  of  nitrogen,  therefore  there 
must  be  three  times  as  much  hydrogen  as  nitrogen 
in  ammonia  gas.  (4)  When  electric  sparks  are 
passed  through  ammonia  gas,  it  is  decomposed  into 
nitrogen  and  hydrogen.  Now  if  oxygen  is  added, 
and  an  electric  spark  passed  through  the  mixture,  the 
oxygen  and  hydrogen  combine.  The  volume  of  the 
remaining  nitrogen  is  one  fourth  of  the  mixture  of 
nitrogen  and  hydrogen,  hence  the  hydrogen  must 
have  been  three  fourths ;  that  is,  the  volume  of 
FIG.  22.— Appa-  hydrogen  in  the  original  volume  ammonia  was  three 
ratus  for  determin-  times  tiiat  of  the  nitrogen.  (5)  The  gravimetric 
iner the  composition  ...  r  .  c  ,  ,  .,.  . 

jr  •  composition  or  ammonia  gas  is  found  by  oxidizing 

it,  and  weighing  the  water  and  nitrogen,  which  are 
the  only  products.  The  result  shows  that  fourteen  parts  of  nitrogen 
combine  with  three  parts  of  hydrogen.  (6)  The  vapor  density  has 
been  found  to  be  8.5.  These  facts  require  NH3  as  the  simplest  formula 
for  ammonia  and  17  as  its  molecular  weight.  Independent  experiments 
verify  this  molecular  weight. 

NITRIC   ACID. 

Nitric  Acid  is  one  of  the  most  useful  compounds  of 
nitrogen.     It  was  known  to  the  alchemists,  who  used  it 


Compounds  of  Nitrogen.  155 

to  prepare  a  mixture  which  dissolves  gold.      Nitric  acid 
is  used  in  the  preparation  of  many  nitrogen  compounds. 

Formation  of  Nitric  Acid.  —  When  moist  animal  or 
vegetable  matter  containing  nitrogen  decays  in  the  presence 
of  an  alkali,  nitric  acid  is  formed ;  it  is  neutralized  at  once 
by  the  alkali,  so  nitrates  —  salts  of  nitric  acid  —  are  the 
final  products.  This  chemical  change  is  known  as  nitri- 
fication, and  it  is  caused,  or  largely  influenced,  by  minute 
living  organisms  called  bacteria.  Nitrification  is  constantly 
going  on  in  the  soil  and  is  an  exceedingly  helpful  process, 
since  it  transforms  harmful  waste  matter  into  valuable 
plant  food. 

As  a  result  of  nitrification,  there  are  vast  deposits  of  nitrates,  espe- 
cially in  desert  regions  and  tropical  countries.  For  example,  potassium 
nitrate  (KNO3)  is  found  in  the  soils  near  large  cities  in  India,  Persia, 
and  Egypt. 

Nitric  acid  is  formed  in  small  quantities  when  electric 
sparks  are  passed  through  moist  air.  Hence  nitric  acid  or 
its  salts  can  be  detected  in  the  atmosphere  after  a  thunder- 
storm. 

This  chemical  change  is  now  being  applied  on  a  large  scale  at  Ni- 
agara Falls.  Electric  sparks  are  passed  through  confined  air  and  the 
products  are  forced  into  a  tower.  Here  they  are  absorbed  in  water  or 
in  a  solution  of  lime ;  thereby  forming  nitric  acid  or  calcium  nitrate. 
The  latter  is  converted  into  sodium  nitrate  (see  below). 

Preparation.  —  Nitric  acid  is  prepared  in  the  laboratory 
by  heating  concentrated  sulphuric  acid  with  a  nitrate,  usu- 
ally sodium  or  potassium  nitrate.  About  equal  weights  of 
nitrate  and  acid  are  put  into  a  glass  retort  and  gently 
heated.  The  nitric  acid  distils  into  a  receiver,  which  is 
kept  cool  by  running  water,  ice,  or  moist  paper.  The 


Descriptive  Chemistry. 


chemical  change  at  a  low  temperature  is  represented  by 
the  equation  — 

NaNO3     +      H2SO4      =   HNO3       +     HNaSO4 
Sodium  Nitrate      Sulphuric  Acid      Nitric  Acid      Acid  Sodium  Sulphate 
85  +98  =         63  +  120 

But  if  the  temperature  is  high  and  an  excess  of  the  nitrate 
is  present,  the  equation  is  — 


2NaNO< 


H2SO4      =   2HNO3 


Na2SO4 
170  +98  126        4-          142 

A  high  temperature,  however,   decomposes    part   of   the 
nitric  acid,  hence  excessive  heat  is  usually  avoided. 


FlG.  23.  —  Apparatus  for  the  manufacture  of  nitric  acid. 

Nitric  acid  is  manufactured  on  a  large  scale  by  heating  sodium  nitrate 
and  sulphuric  acid  in  a  large  cast-iron  retort  (A)  connected  with  huge 
glass  or  earthenware  bottles  (Z?,  B,  £),  arranged  as  shown  in  Figure  23  ; 
the  last  bottle  is  connected  with  a  tower  filled  with  coke  over  which 
water  trickles  to  absorb  the  vapors  which  escape  from  the  bottles.  The 
acid  vapors  are  also  often  absorbed  in  earthenware  or  glass  tubes. 

Properties.  —  Pure  nitric  acid  is  a  colorless  liquid,  but 
the  commercial  acid  is  yellow  or  reddish,  due  to  absorbed 
nitrogen  compounds,  chlorine,  or  iron  compounds.  It  de- 
composes slowly  in  the  sunlight  or  when  heated,  and  a 


Compounds  of  Nitrogen.  157 

brownish  gas  may  often  be  seen  in  bottles  of  nitric  acid. 
It  absorbs  water,  and  forms  irritating  fumes  when  exposed 
to  the  air.  The  specific  gravity  of  the  commercial  acid  is 
about  1.42,  and  it  contains  from  60  to  70  per  cent  of  the 
real  acid  (HNO3),  the  rest  being  water. 

If  the  water  is  removed  by  slowly  distilling  the  commercial  acid  with 
concentrated  sulphuric  acid,  the  product  contains  from  94  to  99  per  cent 
of  the  real  acid  and  its  specific  gravity  is  about  1.51.  When  nitric 
acid  is  boiled,  it  loses  either  acid  or  water  until  the  liquid  contains 
approximately  68  per  cent  of  nitric  acid,  and  then  it  continues  to 
boil  unchanged  at  120°  C. 

Nitric  acid  is  very  corrosive.  It  turns  the  skin  a  perma- 
nent yellow  color,  and  may  cause  serious  burns.  Many 
organic  substances  are  turned  yellow  and  sometimes  com- 
pletely decomposed  by  it.  It  parts  readily  with  its  oxygen, 
especially  when  hot,  and  is  therefore  an  energetic  oxidizing 
agent.  Charcoal  burns  brilliantly  in  hot  acid,  while  straw, 
sawdust,  hair,  and  similar  substances  are  charred  and  even 
inflamed  by  it.  Iron  sulphide  heated  with  nitric  acid 
becomes  iron  sulphate,  by  the  addition  of  oxygen,  thus  — 

FeS      +    2O2    =     FeSO4 

Iron  Sulphide      Oxygen       Iron  Sulphate 

Uses  of  Nitric  Acid.  —  Nitric  acid  is  one  of  the  com- 
mon laboratory  acids.  Large  quantities  are  used  in  the 
manufacture  of  nitrates,  dyestuffs,  sulphuric  acid,  nitro- 
glycerine, gun  cotton,  in  the  refining  of  gold  and  silver, 
and  in  etching  copper  plates. 

Composition  of  Nitric  Acid.  —  Although  the  alchemists  knew  and 
valued  nitric  acid,  its  composition  was  a  mystery  until  Lavoisier  showed 
in  1776  that  it  contained  oxygen  and  probably  nitrogen.  Its  exact 
composition  was  determined  by  Cavendish  in  1784-1785,  by  passing 
electric  sparks  through  a  mixture  of  oxygen  and  nitrogen  in  the  pres- 


158 


Descriptive  Chemistry. 


ence  of  water  or  caustic  potash.  The  same  facts  had  been  observed,  but 
not  explained,  by  Priestley.  Many  independent  experiments  show  that 
the  composition  of  nitric  acid  is  expressed  by  the  formula  HNO3. 

(1)  When  electric  sparks  are  passed  through  a  bottle  containing  moist 
air  or  a  solution  of  potassium  hydroxide,  the  water  becomes  acid  to 
litmus  or  the  liquid  will  be  found  to  contain  a  trace  of  potassium  nitrate. 

(2)  Nitric  acid  may  be  reduced  to  ammonia  by  nascent  hydrogen,  thus 
showing  that  the  acid  contains  nitrogen.     (3)  Conversely,  if  a  mixture 
of  ammonia  and  air  is  passed  over  a  mass  of  hot,  porous  platinum, 
nitric  acid  is  formed.     (4)  If  the  acid  is  allowed  to  flow  through  a  hot 
porcelain  or  clay  tube,  oxygen  is  one  of  the  gaseous  products. 

Nitrates.  —  Nitric  acid  is  monobasic  and  forms  a  series 
of  well-defined  salts  called  nitrates.  The  interaction  of 
nitric  acid  and  most  metals  is  exceedingly  vigorous,  and 
for  this  reason,  probably,  the  alchemists  called  the  acid 
aquafortis  —  strong  water.  The  reaction  varies  with  the 
metal,  strength  of  the  acid,  temperature,  and  the  presence 
of  resulting  compounds. 

The  solid  product  of  the  reaction  is  usually  a  nitrate,  though  some 
metals,  such  as  tin  and  antimony,  form  oxides.  The  gaseous  products 
are  usually  oxides  of  nitrogen,  especially  nitric  oxide  (NO),  which, 
however,  quickly  forms  nitrogen  peroxide  (NO2)  in  the  air.  Hydrogen 
is  never  liberated  so  that  it  can  be  collected  ;  probably  it  immediately 
reduces  the  nitric  acid  to  another  compound  of  nitrogen.  Nitrates  are 
also  formed  by  the  action  of  nitric  acid  upon  oxides,  hydroxides,  and 
carbonates,  thus  — 


CuO 
Copper  Oxide 

KOH 

Potassium 
Hydroxide 

Na2CO3 

Sodium 

Carbonate 


2HNO      = 


HN0      - 


Cu(NO,)2 
Copper  Nitrate 

KN03 

Potassium 
Nitrate 


2HNO3     =     2  NaNO3 
Sodium 
Nitrate 


CO 


H20 


H2O 


H2O 


Compounds  of  Nitrogen.  159 

When  nitric  acid  is  poured  upon  copper,  the  liquid  bub- 
bles violently  and  becomes  hot,  dense  fumes  of  a  reddish 
brown  gas  are  given  off,  and  the  liquid  turns  blue  owing 
to  the  dissolved  copper  nitrate.  Other  metals,  such  as 
zinc,  iron,  and  silver,  act  in  a  similar  way,  though  the  nitrate 
is  blue  only  in  the  case  of  copper.  The  usual  equation  for 
the  chemical  change  with  copper  is  — 

3Cu    +    8HNO3  =   3Cu(NO3)2    +     2  NO      +  4  H2O 
Copper  Nitrate       Nitric  Oxide 

When  nitric  oxide  is  exposed  to  the  air,  it  changes  at  once 
into  the  reddish  brown  peroxide,  thus  — 

NO         +         O  NO2 

Nitric  Oxide  Oxygen  Nitrogen  Peroxide 

Nitrates  as  a  rule  are  very  soluble  in  water.  They  be- 
have in  various  ways  when  heated.  Some,  like  sodium 
and  potassium  nitrates,  lose  oxygen  and  pass  into  nitrites  ; 
others,  like  copper  nitrate,  form  an  oxide  of  the  metal,  an 
oxide  of  nitrogen,  and  oxygen ;  and  one,  ammonium 
nitrate,  decomposes  into  water  and  nitrous  oxide  (N2O). 
Since  many  nitrates,  when  heated,  give  up  oxygen,  they 
are  powerful  oxidizing  agents.  Potassium  nitrate  dropped 
on  hot  charcoal  burns  the  charcoal  vigorously  and  rapidly. 
This  kind  of  chemical  action  is  called  deflagration. 

The  Test  for  Nitrates  (and  of  course  for  nitric  acid)  is  as  follows : 
Add  to  the  solution  of  the  nitrate  a  little  concentrated  sulphuric  acid, 
and  upon  the  cool  mixture  pour  carefully  a  cold,  dilute  solution  of  fresh 
ferrous  sulphate.  A  brown  layer  is  formed  where  the  two  liquids  meet. 

Nitrous  Acid  (HNO2)  has  never  been  obtained  in  the  free  state,  but 
its  salts  —  the  nitrites  —  are  well  known.  Potassium  nitrite  (KNO2) 
and  sodium  nitrite  (NaNO2)  are  formed  by  removing  the  oxygen  from 
the  corresponding  nitrate  by  heating  gently  or  by  heating  with  lead. 
Nitrites  give  off  brown  fumes  when  treated  with  sulphuric  acid,  thus 


i6o 


Descriptive   Chemistry. 


being  readily  di 
decomposition  o 
amount  in  drin 


shed  from  nitrates.     Nitrites  are  formed  by  the 
'C  matter,  and  the  presence  of  a  relatively  large 
r  indicates  contamination  by  sewage. 


Aqua  Regia  is  an  old  term  which  is  still  applied  to  a 
mixture  of  concentrated  nitric  and  hydrochloric  acids. 
The  expression  means  "royal  water,"  and  indicates  that 
the  mixture  dissolves  gold  and  platinum  —  the  noble  metals. 
Its  solvent  power  depends  mainly  upon  the  free  chlorine 
which  is  produced  in  the  mixture  by  the  oxidizing  action 
of  the  nitric  acid.  The  product  of  the  action  of  aqua 
regia  on  metals  is  always  the  chloride  of  the  metal. 

Oxides  of  Nitrogen.  —  There  are  five  oxides  of  ni- 
trogen :  — 


NAME. 

FORMULA. 

CHARACTERISTIC. 

Nitrous  oxide           

N2O 

Colorless  °"as 

NO 

Colorless  o'as 

Nitrogen  trioxide    
Nitrogen  peroxide  

NA 

NO9 

Blue  liquid 
Brown  gas 

Nitrogen  pentoxide      .     .                • 

N,O, 

White  solid 

j.^2^5 

Only  three  of  these  are  important,  viz.,  nitrous  and  nitric 
oxides,  and  nitrogen  peroxide. 

Nitrous  Oxide  is  one  of  the  numerous  decomposition 
products  of  nitric  acid,  but  it  is  usually  prepared  by  decom- 
posing ammonium  nitrate.  This  salt,  if  gently  heated  in  a 
test  tube  provided  with  a  delivery  tube,  first  melts  and  then 
decomposes  into  water  and  nitrous  oxide ;  the  gas  may  be 
collected  over  warm  water.  The  equation  of  the  chemical 
change  is  — 

NH4NO3       =       N2O 

Ammonium  Nitrate      Nitrous  Oxide 


2H2O 


Compounds  of  Nitrogen.  161 

This  colorless  gas  has  a  sweet  taste  and  a  faint  but  pleas- 
ant odor.  It  is  less  soluble  in  hot  than  in  cold  water.  The 
gas  does  not  burn,  but  it  supports  the  combustion  of  many 
burning  substances,  though  not  so  vigorously  as  oxygen 
does.  Sulphur,  for  example,  will  not  burn  in  nitrous  oxide, 
unless  the  sulphur  is  hot  and  well  ignited  at  first.  The 
most  striking  property  of  nitrous  oxide  is  its  effect  on  the 
human  system.  If  breathed  for  a  short  time,  it  causes 
more  or  less  nervous  excitement,  often  manifested  by 
laughter,  and  on  this  account  the  gas  was  called  "laughing 
gas"  by  Davy.  If  breathed  in  large  quantities,  it  slowly  pro- 
duces unconsciousness  and  insensibility  to  pain.  The  gas 
is  often  used  when  insensibility  is  desired  for  a  short  time, 
as  in  dentistry. 

It  is  easily  liquefied  by  cold  and  pressure,  and  is  often  used  in  this 
form  to  furnish  the  gas  itself  and  to  produce  very  low  temperatures.  It 
is  a  commercial  article  and  is  sold  in  small  iron  cylinders. 

Nitrous  oxide  was  discovered  by  Priestley  in  1776;  but  its  composi- 
tion was  not  explained  until  1799,  when  Davy,  by  an  extensive  study  of 
its  properties,  proved  it  to  be  an  oxide  of  nitrogen.  In  his  enthusiasm 
Davy  wrote  a  friend:  "This  gas  raised  my  pulse  upward  of  twenty 
strokes,  made  me  dance  about  the  laboratory  as  a  madman,  and  has 
kept  my  spirits  in  a  glow  ever  since."  It  is  needless  to  say  that  the 
usual  results  are  more  quieting. 

The  Composition  of  Nitrous  Oxide  is  shown  as  follows :  By  ex- 
oloding  equal  volumes  of  nitrous  oxide  and  hydrogen,  only  nitrogen 
Remains,  and  its  volume  equals  the  original  volume  of  nitrous  oxide. 
The  oxygen  unites  with  the  hydrogen  to  form  water,  and  there  is  just 
enough  oxygen  to  unite  with  a  volume  of  hydrogen  equal  to  the  volume 
of  the  nitrous  oxide.  Therefore,  the  oxygen  in  the  nitrous  oxide  must 
have  been  equal  to  half  the  volume  of  the  nitrogen,  since  oxygen  and 
hydrogen  combine  in  the  ratio  of  one  to  two.  Furthermore,  experiment 
has  shown  that  the  weights  of  equal  volumes  of  nitrous  oxide  and  ni- 
trogen are  in  the  ratio  of  44  to  28.  Therefore,  the  smallest  part  of 
oxygen  united  with  the  nitrogen  must  weigh  16 ;  and  since  the  nitrogen 
weighs  28,  the  formula  must  be  N.,0. 


1 62  Descriptive  Chemistry. 

Nitric  Oxide  has  long  been  known,  since  it  is  the  usual 
gaseous  product  of  the  interaction  of  nitric  acid  and  metals. 
It  is  usually  prepared  by  the  interaction  of  copper  and 
dilute  nitric  acid  (sp.  gr.  1.2).  The  equation  for  the  com- 
plex chemical  change  is  usually  written  thus  — 

3Cu  +  8HNO3  =  2  NO  +  Cu(NO3)2  +  4  H2O 
Copper  Nitric  Acid  Nitric  Oxide  Copper  Nitrate 

The  gas  thus  prepared  is  impure,  and  it  is  customary  to 
use  ferrous  sulphate  and  nitric  acid  as  a  source  of  the 
pure  gas. 

Nitric  oxide  is  a  colorless  gas,  but  upon  exposure  to  the 
air,  it  combines  at  once  with  oxygen,  forming  dense  red- 
dish brown  fumes  of  hydrogen  peroxide.  The  simplest 
equation  for  this  change  is  — 

NO       +       O     =       NO2 

Nitric  Oxide  Nitrogen  Peroxide 

This  property  distinguishes  nitric  oxide  from  all  other 
gases.  It  does  not  burn,  nor  does  it  support  combustion 
unless  the  burning  substance  (e.g.  phosphorus  or  sodium) 
introduced  is  hot  enough  to  decompose  the  gas  into  nitro- 
gen and  oxygen,  and  then,  of  course,  the  liberated  oxygen 
assists  the  combustion. 

The  Composition  of  Nitric  Oxide  is  determined  by  heating  iron  or 
another  metal  in  it.  The  oxygen  of  the  oxide  combines  with  the  iron, 
and  the  nitrogen  is  left  free.  The  resulting  volume  of  nitrogen  is  half 
the  volume  of  the  nitric  oxide  taken.  Hence  nitric  oxide  contains 
equal  volumes  of  nitrogen  and  oxygen.  By  an  independent  experiment 
the  molecular  weight  is  found  to  be  30.  Hence  the  formula  must  be  NO. 

Nitrogen  Peroxide  is  the  reddish  brown  gas  formed  by 
the  direct  combination  of  nitric  oxide  and  oxygen.  Thus  — 

NO        +         O  NO2 

Nitric  Oxide  Nitrogen  Peroxide 


Compounds  of  Nitrogen.  163 

It  is  also  produced  by  heating  certain  nitrates.     Thus  — 

Pb(NO3)2    =     2NO2      +       PbO       +       O 
Lead  Nitrate        Nitric  Oxide        Lead  Oxide        Oxygen 

The  fumes  of  nitrogen  peroxide  always  appear  when  nitric 
acid  and  metals  interact,  but,  as  already  stated,  the  fumes 
are  not  produced  at  first,  being  the  result  of  a  second 
chemical  change  when  the  real  product,  nitric  oxide, 
comes  in  contact  with  oxygen  of  the  air. 

Nitrogen  peroxide  is  poisonous.  It  dissolves  in  water ; 
it  also  dissolves  in  concentrated  nitric  acid,  forming 
fuming  nitric  acid. 

At  very  low  temperatures  nitrogen  peroxide  is  a  colorless  solid.  At 
about  — 10°  C.  it  is  a  yellowish  liquid,  and  as  the  temperature  rises  the 
color  grows  darker,  until  at  22°  C.  the  liquid  boils  and  gives  off  the 
familiar  reddish  brown  gas.  Above  140°  C.  this  gas  begins  to  lose  its 
color,  and  at  600°  C.  the  color  entirely  disappears.  The  density  of  the 
gas  at  low  temperatures  indicates  the  formula  N2O4,  whence  the  name 
nitrogen  tetroxide,  often  used.  But  the  density  at  about  140°  C. 
indicates  the  formula  NO2. 

Nitrogen  Trioxide,  N2O3,  and  Nitrogen  Pentoxide,  N2O5,  are  unstable 
compounds  and  have  no  practical  importance.  They  are  the  anhy- 
drides of  nitrous  and  nitric  acids,  thus  — 

N208  +  H2O  2HNO2 

Nitrogen  Trioxide  Nitrous  Acid 

N205  +  H2O  2HNO? 

Nitrogen  Pentoxide  Nitric  Acid 

EXERCISES. 

1.  Name  several  sources  of  ammonia  gas.     How  is  ammonia  gas 
prepared   in   the '  laboratory  ?      Give  the   equation   for  the   reaction. 
State  its  important  properties. 

2.  What  is  ammonium  hydroxide  ?     How  is  it  prepared  on  a  large 
scale  ?     Summarize  its  properties.     What  are  its  uses  ? 

3.  What  is  the  meaning  and   significance   of  (a)    volatile   alkali, 


164  Descriptive  Chemistry. 

($)  anhydrous    ammonia,  (V)    spirits    of  hartshorn,  (d)    sal  volatile, 
(^)  muriate  of  ammonia,  (/")  sal  ammoniac,  (g)  aqua  for'tis  ? 

4.  Why  is  NH3  the  formula  of  ammonia  gas  ? 

5.  Give  several  tests  for  (a)  ammonia,  and  (V)  nitric  acid. 

6.  What  different  meanings  may  the  word  ammonia  have  ?     What 
is   ammoniacal  liquor?     Gas  liquor?     Aqua  ammonia?     Ammonium 
hydrate  ?     Ammonia  of  commerce  ?     Ammonia  water  ? 

7.  How  is  ammonia  gas  liquefied  ?     Describe  the  manufacture  of 
ice  by  liquid  ammonia. 

8.  Develop  the  topics  :  (a)  ammonium  is  a  radical ;   (£)  nitric  acid 
is  an  oxidizing  agent ;  (<:)  nitrates  are  unstable  ;  (d)  fuming  nitric  acid. 

9.  Give  the  formula,  method  of  preparation,  properties,  and  uses  of 
(a)  ammonium  chloride,  (b)  ammonium  nitrate,  (c}  ammonium  sulphate, 
(d)  ammonium  carbonate. 

10.  How  is  nitric  acid  formed  (a)  in  the  soil,  (£)  in  the  air  ?     How 
is  it  prepared  (a)  in  the  laboratory,  (b)  on  a  large  scale  ?     Summarize 
(a)  the  physical  properties  of  nitric  acid,  and  (b}  its  chemical  properties. 
For  what  is  it  used  ? 

1 1 .  What  is  the  formula  of  nitric  acid  ?     Summarize  the  evidence 
of  its  composition. 

12.  What   are   nitrates  ?     How   are   they   formed  ?     What   is   the 
effect  of  heat  upon  («)  potassium  nitrate,  ($)  copper  nitrate,  (c}  am- 
monium nitrate  ?     Give  other  properties  of  nitrates.     What  is  the  test 
for  nitrates  ? 

13.  What   are   nitrites  ?     How   are  they  formed  ?     How  are  they 
distinguished  from  nitrates  ? 

14.  What  is  aqua  regia?    For  what  is  it  used  ?     Why  so  called  ? 
What  is  the  chemical  action  of  aqua  regia  on  gold  ?     Upon  what  prop- 
erty of  nitric  acid  does  its  chemical  action  depend  ? 

15.  Give  the   names  and  formulas  of  the  five  oxides  of  nitrogen. 
Describe  the  preparation  of  nitrous  oxide.     State  briefly  its  properties. 
For  what  is  it  used  ?     Who  discovered  it  ?     What  did  Davy  call  it  ? 
Why  ?     Summarize  the  evidence  of  the  composition  of  nitrous  oxide. 

1 6.  Describe   the  preparation  of  nitric  oxide.     State  the  equation 
for  the  reaction.     What  are  its  properties  ? 

17.  How   is    nitrogen    peroxide    prepared  ?      State    its   properties. 
How   is   it  readily  distinguished  from   all  other  oxides   of  nitrogen  ? 
What  two  formulas  have  been  given  to  nitrogen  peroxide  ?     Why  ? 

1 8.  What  is  (#)  nitric  oxide,  (b)  nitrous  oxide,  (c)  nitrogen  per- 


Compounds  of  Nitrogen.  165 

oxide,    (//)    nitrogen   tetroxide,    (c)    nitrogen    trioxide,    (d)    nitrogen 
monoxide,  (e)  nitrogen  pentoxide  ? 

19.  State  the  equation  for  the  preparation  of  (a)  nitric  acid  at  a 
low  temperature,  (£)  nitric  acid  at  a  high  temperature,  (V)  ammonium 
chloride,  (d)  ammonium  hydroxide  from  water  and  ammonia,  (d}  ni- 
trous oxide,  (e)  nitrogen  peroxide,  (/")  copper  nitrate. 

20.  Define  and  illustrate  («)  sublimation,  ($)  sublimate,  (V)  nitrifi- 
cation, (d}  deflagration,  (>)  nitrate,  (/")  ammonium  compound. 

21.  What  is  the  valence  of  nitrogen  in  ammonia  gas  ?     In  ammo- 
nium ?     In  ammonium  hydroxide  ? 

22.  (a)  Why  are  there  no  acid  nitrates  ?     (£)  What  is  the  valence 
of  nitrogen  in  nitric  acid,  copper  nitrate,  nitrous  oxide,  nitric  oxide, 
nitrogen  peroxide,  nitrogen  trioxide,  nitrogen  pentoxide  ? 


PROBLEMS. 

1.  How  many  grams  of  ammonia  gas  can  be  obtained  from  2140 
gm.  of  ammonium  chloride  by  heating  with  lime  ? 

2.  Calculate  the  percentage  composition  of  (a)  ammonium  chloride, 
(£)  ammonium  hydroxide,  (Y)    ammonium    sulphate,  (d}   ammonium 
nitrate. 

3.  Calculate  the  simplest  formula  of  the  compounds  having  the  per- 
centage composition  (a)  N  =  82.35,  H  =  17.64;    and  (£)  N  =  26.17, 
Cl  =  66.35,  H  =  7.48. 

4.  Calculate  the  percentage  composition  of  (a)  nitric  acid,  (£)  po- 
tassium nitrate  (KNO3),  (V)  sodium  nitrate. 

5 .  How  many  grams  of  nitric  acid  can  be  obtained  by  heating  a 
kilogram  of  sodium  nitrate  with  sulphuric  acid  at  a  low  temperature  ? 

6.  If  the   specific   gravity   of  a   sample   of  nitric   acid   is    1.522, 
(a)  what  will  100  cc.  weigh,  and  (b)  what  volume  must  be  taken  to 
weigh  100  grams  ? 

7.  Calculate  the    simplest  formula  of  the  substances    having   the 
composition    (a)   O  =  76.19,    H  =  1.58,    N  =  22.22;     (£)    N  =13.86, 
K  =  38.61,  O  =  47.52. 


CHAPTER   XIII. 

PROPERTIES  OF  GASES  —  GAY-LUSSAC'S  LAW  OF  GAS 
VOLUMES  —  AVOGADRO'S  HYPOTHESIS  —  VAPOR  DEN- 
SITY AND  MOLECULAR  WEIGHT  —  MOLECULAR  WEIGHTS 
AND  ATOMIC  WEIGHTS  —  MOLECULAR  FORMULA  — MO- 
LECULAR EQUATIONS  —  VALENCE. 

Properties  of  Gases.  —  Extensive  study  of  gases  shows 
that  they  all  conform  to  simple  laws.  Thus  we  have 
already  seen  that  they  behave  uniformly  with  changes  of 
pressure  (Boyle's  law)  and  with  changes  of  temperature 
(Charles's  law).  Other  simple  relations  prevail. 

Gay-Lussac's  Law.  —  Gases  combine  by  volume  in 
simple  ratios.  Experiment  has  revealed  the  following 
facts  about  the  — 

COMBINATION  OF  GASES  BY  VOLUME. 


VOLUMES  OF  COMPONENTS. 

VOLUMES  OF  PRODUCTS. 

2  vol.  hydrogen 
I  vol.  oxygen 

2  vol.  water  vapor 

I  vol.  chlorine 
i  vol.  hydrogen 

2  vol.  hydrochloric  acid  gas 

3  vol.  hydrogen 
i  vol.  nitrogen 

2  vol.  ammonia  gas 

2  vol.  nitrogen 
i  vol.  oxygen 

2  vol.  nitrous  oxide  gas 

2  vol.  nitrogen 
3  vol.  oxygen 

2  vol.  nitrogen  trioxide  gas 

1 66 


Avogadro's  Hypothesis.  167 

Additional  illustrations  will  be  given  in  later  chapters.  The 
simple  ratio  which  exists  between  the  gas  volumes,  whether 
components  or  products,  has  been  found  to  be  true  of  all 
gases.  The  law  was  pointed  out  in  1808  by  Gay-Lussac, 
who  stated  the  relation  substantially  as  follows:  — 

Gases  combine  in  volumes  which  bear  a  simple  ratio  to 
each  other  and  to  that  of  the  product. 

By  "  a  simple  ratio  "  we  mean  one  made  up  of  small 
whole  numbers.  As  a  rule,  the  product  occupies  two  unit 
volumes. 

""  Avogadro's  Hypothesis.  —  In  1811  an  Italian  physi- 
cist proposed  an  hypothesis  to  account  for  the  similar 
behavior  of  gases.  At  that  time  the  properties  of  gases 
were  not  generally  known,  and  the  views  of  Avogadro 
were  overlooked  until  about  1860.  Since  then  the  hypo- 
thesis has  been  helpful  in  explaining  many  facts,  and  it 
is  generally  accepted  by  chemists  as  a  very  probable 
assumption.  It  may  be  stated  thus:  — 

There  is  an  equal  number  of  molecules  in  equal  vohimes 
of  all  gases  at  the  same  temperature  and  pressure. 

This  statement  cannot  be  proved  directly  by  experiment, 
but  there  is  much  physical,  chemical,  and  mathematical 
evidence  in  harmony  with  it. 

According  to  Avogadro's  hypothesis  a  liter  of  hydrogen 
and  a  liter  of  oxygen  at  the  same  temperature  and  pres- 
sure contain  the  same  number  of  molecules,  though  we 
do  not  know  how  many.  Suppose,  however,  that  each 
liter  contained  1000  molecules.  A  liter  of  hydrogen 
weighs  0.0896  gm.  and  a  liter  of  oxygen  at  the  same  tem- 
perature and  pressure  weighs  1.43  gm.  But  0.0896  and 
1.43  are  in  the  same  ratio  as  i  and  16.  Therefore,  since 


1 68  Descriptive  Chemistry. 

a  thousand  molecules  of  oxygen  weighs  16  times  more  than 
a  thousand  molecules  of  hydrogen,  a  single  molecule  of 
oxygen  must  weigh  16  times  more  than  a  single  molecule 
of  hydrogen.  Therefore,  in  general,  in  order  to  find  how 
much  heavier  any  gaseous  molecule  is  than  a  hydrogen 
molecule,  it  is  only  necessary  to  compare  the  weights  of 
equal  volumes  of  hydrogen  and  the  gas  under  examination. 

An  application  of  Avogadro's  hypothesis  is  made  in  course  of  the 
following  argument,  which  proves  that  a  molecule  of  hydrogen  consists 
of  two  atoms  :  — 

One  volume  of  hydrogen  combines  with  one  volume  of  chlorine  to 
form  two  volumes  of  hydrochloric  acid  gas.  Suppose  the  volume  of 
hydrogen  contained  100  molecules.  Then,  according  to  Avogadro's 
hypothesis,  the  equal  volume  of  chlorine  will  contain  100  molecules, 
while  the  two  volumes  of  the  product  will  contain  200  molecules  of 
hydrochloric  acid  gas.  That  is  — 

100  molecules  of  Hydrogen  +  100  molecules  of  Chlorine 
=  200  molecules  of  Hydrochloric  Acid  Gas. 

Now  every  molecule  of  hydrochloric  acid  gas  contains  at  least  one  atom 
each  of  hydrogen  and  chlorine,  and  the  200  molecules  must  contain 
200  atoms  each  of  chlorine  and  hydrogen.  Therefore  each  molecule  of 
hydrogen  and  of  chlorine  must  be  divisible  into  two  atoms,  since  the 
100  hydrogen  and  the  100  chlorine  molecules  provide  the  200  hydrogen 
atoms  and  the  200  chlorine  atoms  in  the  200  molecules  of  hydrochloric 
acid  gas.  Similar  reasoning  leads  to  the  conclusion  that  the  molecules 
of  oxygen,  nitrogen,  and  most  elementary  gases  consist  of  two  atoms. 

Vapor  Density  and  Molecular  Weight.  —  It  was  stated 
in  a  previous  chapter  that  a  molecular  weight  is  the  sum 
of  the  weights  of  the  atoms  in  the  molecule.  But  this 
method  of  finding  the  molecular  weight  is  useless,  unless 
we  first  know  the  formula,  and  in  many  cases  the  formula 
cannot  be  chosen  until  after  the  molecular  weigjit  has  been 
found  by  several  methods.  Hence,  the  determination  of 
molecular  weights  is  an  important  matter.  In  the  case 


Vapor  Density  and   Molecular  Weight.       169 

of  gaseous  or  volatile  elements  and  compounds,  it  is  often 
accomplished  by  finding  the  vapor  density  of  the  substance. 
There  is  a  direct  and  simple  relation  between  molecular 
weight  and  vapor  density.  By  vapor  density  we  mean  the 
ratio  of  the  weight  of  a  gas  to  the  weight  of  an  equal 
volume  of  hydrogen  at  the  same  temperature  and  pressure. 
Thus,  the  vapor  density  of  steam  is  9,  because  experiment 
shows  that  it  weighs  9  times  more  than  an  equal  volume  of 
hydrogen  under  the  same  conditions  of  temperature  and 
pressure.  Therefore  the  molecular  weight  of  steam  is  9 
times  the  molecular"  weight  of  hydrogen.  But  the  molec- 
ular weight  of  hydrogen  is  2,  since  its  molecule  contains 
two  atoms  each  weighing  I.  Therefore,  the  molecular 
weight  of  steam  is  18,  or  twice  the  vapor  density.  The 
general  fact  that  the  molecular  weight  of  a  gaseous  com- 
pound is  twice  its  vapor  density  is  clearly  seen  from  the 
following  table  showing  the  — 

RELATION  BETWEEN  VAPOR  DENSITY  AND  MOLECULAR  WEIGHT. 


GAS. 

VAPOR  DENSITY. 

MOLECULAR  WEIGHT. 

Carbon  dioxide    . 
Ammonia.         .... 

22 
8.5 

44 
17 

Hydrochloric  acid 
Water  vapor  (steam)     . 

18.25 

9 

36.5 

18 

Hence,  a  determination  of  the  vapor  density  of  a  com- 
pound or  an  element  allows  us  to  select  the  correct  molec- 
ular weight  and  assign  the  proper  formula. 

The  vapor  densities  of  the  elements  mercury  and  cadmium  show 
that  the  atom  and  molecule  are  identical,  while  the  vapor  densities  of 
phosphorus  and  arsenic  indicate  that  the  molecule  of  each  consists  of 
four  atoms.  A  molecule  of  oxygen  contains  two  atoms,  but  a  molecule 
of  ozone  contains  three  ;  therefore,  the  formula  of  ozone  is  O3. 


i  jo  Descriptive  Chemistry. 

Other  Methods  of  determining  Molecular  Weights.  —  Some  sub- 
stances cannot  be  vaporized  without  decomposition.  The  molecular 
weights  of  such  substances  cannot,  of  course,  be  found  by  the  vapor 
density  method.  If  a  substance  dissolves  without  decomposition,  its 
molecular  weight  can  be  determined  by  the  boiling-point  or  freezing- 
point  method,  which  was  briefly  described  in  Chapter  X.  The  above 
methods  give  approximate  results.  Exact  molecular  weights  are  found 
by  accurate  quantitative  analysis.  Suppose  we  wished  to  find  the  molec- 
ular weight  of  acetic  acid.  Silver  acetate  is  analyzed  and  found  to  con- 
tain 64.65  per  cent  of  silver ;  the  per  cent  of  the  remaining  elements 
of  the  molecule  must  be  35  35.  The  atomic  weight  of  silver  is  107.93, 
if  the  atomic  weight  of  oxygen  is  16.  Hence,  the  weight  of  the  silver 
acetate  molecule,  except  the  silver,  is  found  by  the  proportion  — 

107.93:  x  :  :  64.65  :  35.35.      x=  59.02. 

Silver  acetate  is  formed  by  replacing  one  atom  of  the  hydrogen  of  the 
acid  by  one  atom  of  silver.  Therefore,  the  weight  of  the  molecule  of 
acetic  acid  is  found  by  adding  to  59.02  the  weight  of  one  atom  of  hydro- 
gen. That  is,  the  exact  molecular  weight  of  acetic  acid  is  60.028 
(i.e.  59.02  +  1.008). 

Determination  of  Atomic  Weights. — The  atomic 
weight  of  an  element,  as  already  stated,  is  a  relative  weight. 
It  is  a  number  expressing  the  relation  of  the  weight 
of  an  atom  of  a  given  element  to  the  weight  of  an  atom 
of  some  element  chosen  as  a  standard.  Thus,  if  we  say 
that  the  atomic  weight  of  nitrogen  is  14,  we  mean  that 
the  relation  between  the  weight  of  the  nitrogen  atom  and 
that  of  the  hydrogen  atom  is  14  to  i,  if  we  adopt  the 
hydrogen  atom  as  the  standard  atom ;  or  we  mean  that 
the  relation  between  the  weight  of  the  nitrogen  atom  and 
that  of  the  oxygen  atom  is  14  to  16,  if  we  adopt  the  oxygen 
atom  as  the  standard.  The  approximate  atomic  weights 
are  usually  expressed  in  round  numbers,  and  do  not 
vary  much  with  the  standard.  Wherever  exact  atomic 
weights  are  used  in  this  book,  the  oxygen  standard  is  the 
basis. 


Determination  of  Atomic  Weights.  171 


In  Chapter  IX  it  was  stated  that  the  determination  and 
selection  of  atomic  weights  are  based  on  several  principles. 
This  subject  can  now  be  appropriately  considered. 

One  method  of  selecting  the  atomic  weight  is  illustrated 
by  the  case  of  chlorine,  which  has  the  atomic  weight  35.5. 
The  molecular  weights  of  several  chlorine  compounds  are 
found  by  the  vapor  density  method.  The  compounds  are 
analyzed  to  find  the  number  of  grams  of  chlorine  in  the 
number  of  grams  of  the  compound  equal  to  the  determined 
molecular  weight.  And  the  highest  common  factor  of 
these  weights  of  chlorine  is  taken  as  the  atomic  weight  of 
the  element.  A  concise  view  of  the  method  is  shown  in 
the  following  — 

TABLE  OF  CHLORINE  COMPOUNDS. 


COMPOUND. 

MOLECULAR 
WEIGHT. 

WEIGHT  OF 
CHLORINE. 

H.  C.  F. 

Hydrochloric  acid  .... 
Chlorine  peroxide   .... 
Cyanogen  chloride  .... 
Chlorine  sas 

36.5 
67.5 
6l.S 
71 

35-5 
35-5 
35-5 

71 

i  x  35-5 
i  x  35-5 
i  x  35.5 

2  X   3C   C 

Chlorine  monoxide      .     .     . 
Phosphorus  trichloride     .     . 
Chloroform     

8? 
137.5 
I  IQ.1 

71 
106.5 

io6.cr 

•*•  JJ'J 
2  X  35.5 

3  x  35.5 
?  x  "K.c( 

Carbon  tetrachloride    .     .     . 

170 

142 

4  x  35.5 

Thirty-five  and  five  tenths  is  therefore   selected   as   the 
approximate  atomic  weight  of  chlorine. 

Atomic  weights  can  also  be  determined  by  analysis  if 
we  know  the  proportion  in  which  the  atoms  combine  to 
form  a  molecule  of  the  compound  analyzed.  Thus,  the 
Belgian  chemist,  Stas,  who  made  masterly  determinations 
of  atomic  weights,  found  that  121.4993  gm.  of  silver 


172  Descriptive   Chemistry. 

chloride  were  formed  by  burning  91.462  gm.  of  silver 
in  chlorine.  He  knew  that  one  atom  of  silver  and  one 
of  chlorine  unite  to  form  silver  chloride  ;  he  also  accepted 
35.453  as  the  atomic  weight  of  chlorine.  Hence,  he  calcu- 
lated the  atomic  weight  of  silver  thus  — 

121.4993-91.462=  30.0373, 

which  is  the  weight  of  the  chlorine  used. 
Therefore  — 

91.462  :  30.0373  : :  x  :  35-453,    *  =  107.95, 

the  atomic  weight  of  silver. 

Approximate  atomic  weights  of  the  solid  elements,  espe- 
cially the  metals,  are  checked  by  applying  the  law  of  spe- 
cific heats.  This  law  was  announced  by  Dulong  and 
Petit  in  1819.  It  is  stated  as  follows :- 

The  product  of  the  specific  heat  and  atomic  weigJit  of  tJie 
solid  elements  is  a  constant  quantity. 

By  Specific  heat  we  mean  the  quantity  of  heat  necessary 
to  raise  the  temperature  of  a  substance  one  degree  com- 
pared with  the  quantity  necessary  to  raise  the  temperature 
of  the  same  weight  of  water  one  degree.  If  the  same 
quantity  of  heat  is  imparted  to  equal  weights  of  water  and 
mercury,  the  temperature  of  the  mercury  will  be  much 
higher  —  about  32  times  higher  than  that  of  the  water. 
That  is,  the  mercury  requires  only  about  ^  as  much  heat 
as  the  water.  In  other  words,  the  specific  heat  of  mercury 
is  ^2>  or  o-°3 r  2-  The  specific  heat  of  other  elements  is  simi- 
larly found. 

The  constant  quantity  found  by  multiplying  the  specific 
heat  by  atomic  weight  is  approximately  6.25.  This  rela- 
tion is  illustrated  by  the  following  — 


Determination  of  Atomic  Weights.  173 

TABLE  OF  SPECIFIC  HEATS. 


ELEMENT. 

SPECIFIC  HEAT. 

ATOMIC  WEIGHT. 

PRODUCT. 

Calcium 

O  I7O 

AQ 

6  8 

Copper 

O  OQC 

6^  6 

6  04. 

Iron    

'•^yj 
O.I  14. 

c6 

638 

Lead  
Potassium    . 
Sodium              .... 

0.031 
0.166 
O.2Q'? 

207 

39 

2"J 

6.41 

6.47 

6  7^ 

Sulphur 

0.178 

M 

57 

Tin 

O  OCC 

I  IQ 

•/ 

6  cj. 

Zinc 

->.W-J} 
O  OQ4. 

6c  A 

6  i  c 

•'•wyif 

^y't 

The  use  of  this  law  in  checking  atomic  weights  may  be 
illustrated  as  follows :  The  specific  heat  of  silver  is  found 
by  experiment  to  be  0.057;  if  6.25  is  divided  by  this  num- 
ber, the  quotient  is  approximately  109.  This  result  agrees 
approximately  with  108  —  the  accepted  atomic  weight  of 
silver.  Again,  the  specific  heat  of  mercury  is  0.0312;  if 
6.25  is  divided  by  this  number,-  the  quotient,  200,  indicates 
that  the  atomic  weight  of  mercury  is  200  —  a  value  obtained 
by  other  methods.  This  law  has  been  of  assistance  in  the 
final  selection  of  the  approximate  atomic  weight  of  several 
elements.  Thus,  the  atomic  weight  of  uranium  was  finally 
accepted  as  about  238  instead  of  119.  Both  values  agreed 
with  analyses,  but  only  the  former  conformed  to  Dulong 
and  Petit's  law. 

The  plan  followed  in  determining  the  atomic  weight  of  zinc  illustrates 
the  methods  actually  used. 

(a)  When  zinc  interacts  with  dilute  hydrochloric  or  sulphuric  acid, 
hydrogen  is  liberated  ;  and  if  a  known  weight  of  zinc  is  used,  the  weight 
of  zinc  needed  to  liberate  I  gm.  of  hydrogen  is  easily  calculated. 
This  number,  as  we  have  already  seen,  is  the  equivalent  of  zinc  (see 
Equivalents,  Chapter  IX).  Now  if  one  atom  of  zinc  replaces  one  atom 


174  Descriptive  Chemistry. 

of  hydrogen,  then  the  atomic  weight  of  zinc  and  the  atomic  weight  of 
hydrogen  will  have  the  same  ratio  as  the  weight  of  zinc  and  the  weight 
of  hydrogen  found  by  experiment.  According  to  experiment  the 
equivalent  of  zinc  is  about  32.5.  This  is  its  relation,  atom  for  atom, 
to  hydrogen,  and,  thus  far,  is  its  atomic  weight. 

(£)  When  zinc  and  hydrochloric  acid  interact,  zinc  chloride  is 
formed.  If  it  is  analyzed,  the  proportion  of  zinc  to  chlorine  is  about 
32.5  to  35.5.  If  the  elements  combine,  atom  for  atom,  the  atomic 
weight  of  zinc  is  32.5  (assuming  that  35.5  is  the  atomic  weight  of 
chlorine). 

(<:)  When  zinc  is  burned  in  air,  zinc  oxide  is  formed.  If  this  com- 
pound is  analyzed,  the  proportion  of  zinc  to  oxygen  is  about  65  to  16. 
If  the  elements  combine  atom  for  atom,  the  atomic  weight  of  zinc  is 
about  65  (assuming  that  16  is  the  atomic  weight  of  oxygen). 

(df)  According  to  these  three  determinations,  the  atomic  weight  of 
zinc  is  32.5  or  65.  We  have  assumed  that  the  elements  unite  atom  for 
atom  in  each  compound.  This  is  an  incorrect  assumption,  because  an 
atom  of  zinc  cannot  have  two  different  weights  —  32.5  and  65.  If  the 
atomic  weight  is  32.5,  zinc  oxide  must  consist  of  one  atom  of  oxygen 
and  two  of  zinc.  But  if  the  atomic  weight  is  65,  zinc  chloride  must 
consist  of  two  atoms  of  chlorine  and  one  of  zinc,  and  two  atoms  of 
hydrogen  must  have  been  replaced  by  one  of  zinc. 

(e)  The  molecular  weight  of  zinc  chloride  is  found  by  the  vapor 
density  method  to  be  about  133.  If  zinc  chloride  consists  of  two 
atoms  of  chlorine  and  one  of  zinc  (weighing  65),  its  molecular  weight 
is  about  136.  In  other  words,  it  is  evident  that  our  assumption  regard- 
ing the  number  of  atoms  in  zinc  chloride  is  highly  probable. 

CO  We  are  not  absolutely  positive,  however,  that  the  zinc  in  a 
molecule  of  zinc  chloride  may  not  be  one  atom  weighing  65,  or  two 
atoms  weighing  32.5  each.  But  the  atomic  weight  of  zinc  determined 
by  applying  the  law  of  specific  heats  is  664  (i.e.  6.25  -:-  0.094).  This 
shows  clearly  that  the  atomic  weight  of  zinc  is  approximately  65. 

Molecular  Formula.  —  In  Chapter  IX  a  method  was 
given  for  finding  the  simplest  formula  of  a  compound,  viz., 
by  dividing  the  percentage  of  each  element  by  its  atomic 
weight.  But  the  simplest  formula  is  not  always  the  mo- 
lecular formula ;  that  is,  it  does  not  always  express  the 
composition  and  number  of  atoms  in  a  molecule  of  the 


Molecular  Equations.  175 

compound  in  the  gaseous  state.  Every  formula,  however, 
is  designed  to  be  a  molecular  formula.  Since  the  molecu- 
lar weight  of  a  compound  is  twice  its  vapor  density,  the 
molecular  formula  can  be  calculated  from  the  simplest 
formula.  Thus,  the  simplest  formula  of  a  compound  of 
carbon  and  hydrogen  was  found  to  be  CH2.  Its  vapor 
density  was  found  to  be  81.4.  Hence  its  molecular  weight 
must  be  162.8,  which  is  nearly  twelve  times  that  corre- 
sponding to  CH2.  Therefore  the  molecular  formula  is 
C12H24.  Molecular  formulas  of  other  compounds  may  be 
similarly  found. 

Molecular  Equations.  —  Equations  which  represent  re- 
actions between  gases  are  sometimes  written  as  molecular 
equations.  Such  equations  represent  changes  as  taking 
place  between  the  smallest  possible  physical  units,  that  is, 
between  molecules.  The  molecular  equation  for  the  for- 
mation of  water  from  hydrogen  and  oxygen  is  — 

2  H2     +     O2     =     2  H2O. 

It  is  read  thus  :  Two  molecules  of  hydrogen  unite  with  one 
molecule  of  oxygen  to  form  two  molecules  of  water.  Since 
most  elementary  gases  consist  of  molecules,  such  an  equa- 
tion is  strictly  correct.  It  should  be  noted,  however,  that 
the  proportions  are  the  same  as  in  the  simpler  form  of  the 
equation.  For  practical  purposes  the  molecular  equation 
is  preferable  only  in  the  case  of  gases. 

Molecular  equations  are  sometimes  called  volume  or  gas  equations, 
because  such  equations  tell  at  a  glance  the  volumes  involved  in  the  re- 
action. Thus-  H2  +  Cl2  =  2HC1 

means  that  one  volume  each  of  hydrogen  and  chlorine  unite  to  form  two 
volumes  of  hydrochloric  acid  gas.  This  equation  is  sometimes  writ- 

ten  —  H2     +       C12     =     2HC1 

I  VOL  I  VOl.  2  VOl. 


176  Descriptive  Chemistry. 

Valence.  — An  examination  of  many  formulas  obtained 
by  the  principles  just  discussed  shows  certain  regularities. 
Take,  for  example,  some  binary  compounds  of  hydrogen. 
They  fall  into  four  groups,  thus  — 

I.  II.  III.  IV. 

HC1  H2O  H3N  H4C 

HBr  H2S  H3P  H4Si 

Obviously,  the  atoms  of  these  elements  differ  in  their 
power  of  combining  with  hydrogen  atoms.  Some  unite 
with  one  atom,  some  with  two  atoms,  and  so  on.  Atoms 
of  other  elements  besides  those  in  the  above  list  differ  in 
their  combining  power.  The  power  of  atoms  of  an  ele- 
ment to  hold  in  combination  a  certain  number  of  other 
atoms  is  called  the  valence  or  quantivalence  of  the 
element.  The  valence  of  hydrogen  is  always  one.  Ele- 
ments which  combine  atom  for  atom  with  one  atom  of 
hydrogen  have  the  valence  one,  and  are  called  univalent 
elements  or  monads ;  sodium  and  potassium  are  always 
univalent,  and  so  is  chlorine  in  hydrochloric  acid.  Ele- 
me'nts  which  combine  with  two  atoms  of  hydrogen  have 
the  valence  two,  and  are  called  bivalent  elements  or 
dyads ;  oxygen,  magnesium,  and  sulphur  are  bivalent 
elements.  So,  also,  some  elements  like  aluminium,  are  tri- 
valent  or  triads ;  others,  like  carbon  and  silicon,  are 
quadrivalent  or  tetrads;  and  some,  like  the  nitrogen  in 
nitric  acid,  are  quinquivalent  or  pentads.  Elements  of 
the  same  valence  combine  with  or  replace  each  other  atom 
for  atom.  Thus,  one  atom  of  sodium  replaces  one  atom 
of  hydrogen  in  hydrochloric  acid ;  and  one  atom  of  oxygen 
combines  with  one  atom  of  magnesium.  Elements  of  dif- 
ferent valence  form  compounds  in  which,  as  a  rule,  the 
number  of  atoms  is  such  that  the  valences  balance,  Thus, 


Valence.  177 

a  dyad  combines  with  two  monads  (as  in  H2O),  a  triad  with 
three  monads  (as  in  NH3),  two  triads  with  three  dyads  (as 
in  A12O3),  one  tetrad  with  two  dyads  (as  in  CS2),  and  so 
on.  Such  compounds,  in  which  the  capacity  for  further 
union  has  ceased,  are  said  to  be  saturated  or  to  have  no 
free  bonds.  Compounds  in  which  the  valence  is  not  bal- 
anced, or  in  which  free  bonds  exist,  are  called  unsaturated 
(see  Ethylene). 

The  valence  of  an  element  is  always  the  same  in  the 
same  compound,  but  it  often  varies.  Thus,  the  valence  of  ni- 
trogen is  one  in  N2O,  two  in  NO,  three  in  N2O3,  four  in  NO2, 
and  five  in  HNO3.  Hydrogen,  as  stated  above,  always 
has  a  valence  of  one  ;  it  is  also  believed  that  the  valence  of 
oxygen  is  always  two.  If  an  element  forms  no  hydrogen 
compound,  its  valence  is  determined  from  compounds  con- 
taining elements  which  are  univalent,  such  as  chlorine, 
bromine,  and  sodium. 

The  valence  of  elements  in  saturated  compounds  of  two 
elements  is  easily  deduced  from  the  formula,  because  in 
such  compounds  the  total  valence  of  all  the  atoms  of  each 
element  must  be  equivalent  Thus  in  the  formula  CaO, 
the  valence  of  calcium  is  two,  because  the  single  atom  of 
calcium  is  combined  with  a  single  atom  of  a  bivalent 
element.  The  valence  of  phosphorus  in  P2O5  is  five,  be- 
cause the  two  atoms  furnish  a  total  valence  of  ten,  which 
is  required  by  the  five  atoms  of  the  bivalent  element 
oxygen.  In  CH4  the  valence  of  carbon  is  four,  because  the 
single  atom  is  combined  with  four  atoms  of  hydrogen. 

Radicals  have  a  valence,  since  in  chemical  changes  they 
act  like  atoms.  The  valence  of  ammonium  (NH4)  is  one, 
and  of  hydroxyl  (OH )  is  one.  Thus,  NH4C1  is  the  formula 
of  ammonium  chloride,  NaOH  of  sodium  hydroxide,  but 
Ca(OH)3  of  calcium  hydroxide. 

' 


1  7  8  Descriptive  Chemistry. 

The  valence  of  elements  in  unsaturated  compounds  can- 
not be  told  by  mere  inspection  ;  a  knowledge  of  the  prop- 
erties of  the  compound  is  necessary.  So  also  the  valence 
of  some  elements  in  compounds  containing  three  or  more 
elements  is  not  readily  told  from  the  formulas^  some 
knowledge  of  the  methods  of  formation,  relations  to  other 
compounds,  and  general  properties  is  needed.  A  discus- 
sion of  these  principles  is  beyond  the  scope  of  this  book. 
However,  in  the  case  of  most  acids,  bases,  and  salts,  an 
arbitrary  rule  may  be  cited.  In  these  compounds  the  total 
valence  of  the  oxygen  atoms  balances  the  total  valence  of 
the  other  elements.  Thus,  in  nitric  acid,  HNO3,  the  va- 
lence of  nitrogen  is  nve,  while  in  nitrous  acid,  HNO2,  it  is 
three. 

Some  chemists  prefer  to  regard  valence  as  the  quotient  obtained  by 
dividing  the  atomic  weight  by  the  equivalent  weight.  For  example, 
the  valence  of  oxygen  is  2  —  the  quotient  of  16  -4-  8.  Such  a  view  is 
not  inconsistent  with  the  one  generally  held,  because  valence  is  the 
direct  outcome  of.  composition. 

The  valence  of  elements  may  be  represented  in  several  ways,  e.g. 

H',  H  —  ,  —  O  —  ,  O  =  ,  N  —  .      Sometimes  formulas  are  written  to 

show  the  valence,  e.g.  —  /  H 

Hydrochloric  acid,  H  -  Cl,  Water,  H  -  O  -  H,  Ammonia,  N  -  H. 


Such  formulas  are  called  structural  or  graphic  formulas  to  distinguish 
them  from  the  ordinary  or  empirical  formulas.  Structural  formulas 
are  not  intended  to  show  how  the  atoms  are  arranged  in  space.  We 
know  very  little  about  the  space  relations  of  atoms.  They  simply  indi- 
cate certain  relations  not  shown  by  the  empirical  formulas.  They  are 
especially  helpful  in  organic  chemistry  (see  Chapter  XXXI). 

EXERCISES. 

i  .    Review  (a)  Boyle's  law,  and  (ft)  Charles's  law. 

2.  State  and  illustrate  Gay-Lussac's  law. 

3.  Give  a  brief  account  of  (a)  Gay-Lussac?  (£)  Avogadro,  (c)  Stas, 


Exercises. 


179 


4.  State  and  illustrate  Avogadro's  hypothesis. 

5.  What  is  the  relation  of  the  molecular  weight  of  a  gas  to  (a)  the 
molecular  and  (£)  the  atomic  weight  of  hydrogen  ? 

6.  (a)  State  the  argument  proving  that  a  molecule  of  hydrogen  con- 
sists of  two  atoms,     (b}  Apply  the  same  argument  to  oxygen. 

7.  What  is  the  relation  between  molecular  weight  and  vapor  den- 
sity ?      Illustrate  your  answer.      What  application  is   made  of  this 
relation  ? 

8.  Why  is  the  formula  of  water  H2O  and  not  HO  or  H2O2  ? 

9.  Why  is  the  formula  of  ozone  O3  ? 

10.  (a)  How   are    molecular   weights    determined  ?    (£)  How  are 
atomic  weights  found  from  molecular  weights  ? 

11.  Illustrate  the  method  of  determining  atomic  weights  by  chemical 
analysis. 

12.  What  is  a  molecular  formula  ?     What  is  the  molecular  formula 
of  oxygen,  nitrogen,  chlorine,  and   hydrogen  ?     How  is  a  molecular 
formula  determined  ?     Illustrate  your  answer. 

13.  What  is  a  molecular  equation  ?     Give  two  illustrations.     How 
does  it  differ  from  an  ordinary  chemical  equation  ?     Of  what  use  are 
such  equations  ? 

14.  Define    («)  valence,   (b}   monad,  dyad,    triad,  tetrad,  pentad, 
(c)    univalent    element,  bivalent    element,   (d)    saturated    compound, 
(e)  unsaturated  compound. 

15.  What  is  the  valence  of  hydrogen  ?    Why  ?    Of  oxygen  ?  Why  ? 
How  may  valence  be  found  by  inspecting  a  binary  formula  ?     What  is 
the  valence  of  NH4  and  OH  ? 

1 6.  Illustrate  the  ways  valence  may  be  represented. 

17.  Distinguish  between  structural  and  empirical  formulas. 

1 8.  What  is  the  valence  of  sodium  in  (a}  sodium  chloride,  (b)  so- 
dium nitrate  (NaNO3),  (c}  sodium   sulphate    (Na2SO4),  (d)    sodium 
hydroxide  (NaOH)  ? 

19.  What  is  the  valence  of  sulphur  in  (a)  snlphur  dioxide  (SO2), 
(b)  sulphur  trioxide  (SO3),  (c}  hydrogen  sulphide  (H2S),  (d)  sulphuric 
acid,  (e)  copper  sulphate  (CuSO4)  ?     (Suggestion.  —  In  oxygen  acids, 
the  oxygen  valence  balances  the   sum   of  the  valence   of  the   other 
elements.) 

20.  What   is  the  valence  of  (#)    aluminium   in   aluminium   oxide 
(Al2Oo),  (b}  carbon  in  carbon  tetrachloride  (CC14),  (c)  phosphorus  in 
phosphorus  pentoxide  (P2O.)  ? 


I 


180  Descriptive  Chemistry. 

21.  What  is  the  valence  of  (a}  silver  and  chlorine  in  silver  chloride 
(AgCl),  (£)  calcium  and  chlorine  in  calcium  chloride  (CaCl2),  (<;)  oxy- 
gen in  water,  (d)  oxygen  and  calcium  in  calcium  oxide  or  lime  (CaO)  ? 

PROBLEMS. 

1.  The  vapor  densities  of  certain  gases  is  as  follows :   (#)  hydro- 
chloric acid  18.25,  (b}  chlorine  35.5,  (c)  ammonia  8.5,  (d)  nitrogen  14, 
0)  steam  9.     Calculate  the  molecular  weight  of  each. 

2.  Calculate  the  simplest  formula  of  the  compounds  which  have 
the  indicated  composition:   (a}  N  =  82.353,  H  =  17.647;   (£)  O  =  30, 
Fe  (iron)  =70;   (c)  H  =  i,  C  =  12,  K  (potassium)  =39,  O  =  48. 

3.  A    liter    of    sulphurous    oxide   gas  (SO2)  weighs  2.8672    gm. 
What  is  the  molecular  weight  of  this  compound  ? 

4.  If  1500  cc.  of  carbon  monoxide  gas  (CO)  weigh    1.8816  gm., 
what  is  the  molecular  weight  of  the  compound  ? 

5.  Calculate  the  molecular  formula  of  the  compounds  corresponding 
to  the  following  data:    (a)   C  =  73.8,  H  =  8.7,  N  =  17.1,  vapor  density 
=  80.2;   (£)  C=92.3,    H  =  7.7,  vapor    density  =38.8 ;    (c)  C  =  39.9, 
H  =  6.7,  O  =53.4,  vapor  density  =  30.5. 

6.  What  volumes  of  factors  and  products  are  represented  by  the 
equations  (a)  H2  +  C12  =  2  HC1,  (£)  2  H,  +  O2  =  2  H,O,  (c)  3  H,  + 
N2  =  2  NH3,  (d)  N2  +  O2  =  2  NO,  (e)  2  NO  +  O,  =  2  NO,  ? 

7.  If  20  1.  of  hydrogen  are  allowed  to  interact  with  10  1.  of  chlo- 
rine, (a)  how  many  liters  of  hydrochloric  acid  gas  are  produced,  and 
(£)  which  gas  and  how  much  remains  ? 

8.  How  many  liters  of  hydrogen  gas  can  be  obtained  from  4  1. 
of  hydrochloric  acid  gas  ? 

9.  If  91.462  gm.  of  silver,  when  heated  in  chlorine,  yield  121.4993  gm. 
of  silver  chloride,  what  is  the  atomic  weight  of  chlorine  ?     (Assume 
Ag  =  108.) 

10.  How  many  liters  of  the  component  gases  can  be  obtained  by 
the  decomposition  of  6  1.  of  ammonia  gas  ? 

11.  Find  the  simplest  formulas  of  the  substances  having  the  follow- 
ing composition  :   (a)  H  =  1.58,  N  =  22.22,  O  =  76.19 ;   (^)  O  =  47.52, 
N  =  13.86,  K  =  38.61. 

12.  A   certain  weight   of  copper  oxide,  when  heated  in  a  current 
of  hydrogen,  lost   59.789   gm.   of  oxygen  and  formed  67.282  gm.  of 
water.      (a)    If  O  =  16,   what    is    the    atomic   weight   of  hydrogen  ? 
(£)   If  H  =  i,  what  is  the  atomic  weight  of  oxygen  ? 


CHAPTER   XIV. 
CARBON  AND   ITS  OXIDES  —  CYANOGEN. 

Occurrence  of  Carbon.  —  Uncombined  carbon  is  found 
pure  in  nature  as  diamond  and  graphite ;  in  a  more  or  less 
impure  state  it  occurs  as  coal  and  similar  substances, 
which  are  included  in  the  term  amorphous  carbon.  Car- 
bon forms  a  vast  number  of  compounds,  natural  and 
artificial.  Combined  with  hydrogen  and  oxygen,  and 
occasionally  with  nitrogen  also,  it  is  an  essential  constitu- 
ent of  plants  and  animals.  Meat,  starch,  fat,  sugar,  wood, 
cotton,  paper,  soap,  wool,  wax,  flour,  albumen,  and  bone 
contain  carbon.  It  is  also  a  component  of  carbon  dioxide 
and  of  carbonates,  such  as  limestone,  chalk,  and  marble. 
Illuminating  gases,  kerosene  and  other  products  of  petro- 
leum, turpentine,  alcohol,  chloroform,  ether,  and  similar 
liquids  are  compounds  of  carbon.  It  is  estimated  that 
0.22  per  cent  of  the  weight  of  the  earth's  crust  is  carbon. 

Diamond  is  pure  crystallized  carbon.  It  is  found  in 
only  a  few  places  in  the  earth.  When  taken  from  the 
mine,  diamonds  are  rough-looking  stones ;  some  are  crystals, 
some  are  rounded  like  peas,  and  many  are  irregular ;  they 
must  be  cut  and  polished  to  bring  out  the  luster  and  make 
them  sparkle  (Fig.  24).  The  highly  prized  diamonds  are 
colorless  and  without  a  flaw,  and  are  said  to  be  "of  the 
first  water  "  ;  yellow  ones  from  South  Africa  are  common, 
and  occasionally  a  blue,  pink,  red,  or  green  one  is  found ; 
a  very  impure  variety  is  black. 

181 


182 


Descriptive  Chemistry. 


The  diamond  is  insoluble  in  all  liquids  at  the  ordinary 
temperature,  has  the  high  specific  gravity  of  3.5,  and  is  the 
hardest  known  substance. 

It  is  brittle  and  may  be  shattered  by  a  blow  with  a 
hammer. 


Crystal. 


Rough. 
FIG.  24.  —  Diamonds. 


Cut. 


Diamonds  have  always  been  prized  as  gems  on  account  of  their 
beauty,  rarity,  and  permanency.  Besides  being  worn  as  jewels,  they 
are  used  to  cut  glass,  and  the  powder  and  splinters  (known  as  bort) 
are  used  to  grind  and  polish  diamonds  and  other  hard  gems.  The  im- 
pure variety  which  comes  from  Brazil,  and  is  called  carbonado,  is  set  into 
the  end  of  the  "  diamond  drill,'1  which  is  used  extensively  for  boring 
artesian  wells  and  drilling  hard  rocks. 

The  diamond  was  formerly  found  in  gravel  deposits  in  India,  and  in 
later  years  in  Brazil.  Since  1867,  however,  about  95  per  cent  of  the  dia- 
monds of  commerce  have  come  from  South  Africa.  They  occur  in  a 
bluish  volcanic  rock  along  the  Vaal  River,  and  especially  near  Kimberley. 
Over  eight  tons  of  diamonds  have  been  found  in  South  Africa  in  the 
last  twenty-five  years  ! 

The  successive  investigations  of  Lavoisier,  Dumas,  and  Davy,  ex- 
tending from  1772  to  1814,  showed  that  diamond  is  carbon,  for  when 
pure  diamond  was  burned  in  oxygen,  the  only  product  was  carbon 

dioxide.  This  result,  which  ad- 
mits of  no  doubt,  has  been  verified 
by  many  famous  investigators. 
Diamonds  have  been  made  by 
Moissan.  He  dissolved  pure  char- 
coal in  melted  iron,  and  poured  the 
molten  mass  into  water.  The  sur- 
face was  so  suddenly  cooled  that  a  tremendous  pressure  was  exerted 


FlG.  25. —  Artificial  diamonds  (enlarged) 
prepared  by  Moissan. 


Carbon  and  its  Oxides.  183 

by  the  expanding  iron  inside  the  crust.  This  pressure  caused  the  cool- 
ing carbon  to  crystallize  into  diamond.  The  crystals  were  very  small, 
most  of  them  were  black,  a  few  were  white,  but  all  had  the  properties 
of  the  diamond  (Fig.  25). 

Large  diamonds  have  a  fascinating  history,  since  most  of  them  have 
passed  through  many  hands  before  finding  a  place  among  royal  jewels. 
The  largest  is  the  Orloff,  which  weighs  194!  carats,  and  is  in  the  scepter 
of  the  Czar  of  Russia.1  The  Kohinoor,  which  now  weighs  about  106 
carats,  is  one  of  the  crown  jewels  of  England. 

Graphite  is  a  soft,  black,  shiny  solid,  which  is  smooth 
and  soapy  to  the  touch.  Pure  graphite  is  carbon.  It  occurs 
native  in  large  quantities  and  in  many  places.  One  va- 
riety is  found  in  abundance  at  Ticonderoga,  New  York. 
Other  famous  localities  are  Ceylon,  eastern  Siberia,  Bava- 
ria, and  Italy.  Sometimes  crystals  and  grains  are  found, 
but  it  usually  occurs  in  flaky  masses  or  slabs.  Unlike 
diamond,  graphite  is  a  good  conductor  of  electricity  and  is 
often  used  to  coat  moulds  in  electrotyping.  It  is  so  soft 
that  it  blackens  the  fingers  and  leaves  a  black  mark  on 
paper  when  drawn  across  it.  This  property  is  indicated 
by  the  name  graphite,  which  is  derived  from  a  Greek  word 
(grap/iein)  meaning  to  write.  It  resembles  diamond  in  its 
insolubility  in  liquids  at  the  ordinary  temperature.  Its 
specific  gravity  is  2.2,  being  considerably  lighter  than  dia- 
mond. It  produces  only  carbon  dioxide  when  burned  in 
oxygen ;  but  unlike  diamond,  it  turns  into  carbon  dioxide 
by  heating  to  a  very  high  temperature  in  the  air.  Graphite 
was  once  supposed  to  contain  lead,  and  rs  even  now  often 
incorrectly  called  "  black  lead  "  and  plumbago.  It  is  used 
to  make  stove  polish  and  protective  paints,  as  a  lubricant 
where  oil  cannot  be  used,  as'  the  principal  ingredient  of 

1  A  carat  equals  3J  Troy  grains  (or  0.205  gm.).  The  term  is  derived  from  the 
carob  bean,  which  was  used  for  ages  by  the  diamond  merchants  of  India  as  a 
small  weight. 


184 


Descriptive  Chemistry. 


graphite  crucibles,  in  which  metals  are  often  melted,  and 
in  making  electrodes  for  the  huge  electric  furnaces. 

Immense  quantities  of  graphite  are  consumed  in  the  manufacture  of 
lead  pencils.  The  graphite  is  washed  free  from  impurities,  ground  to  a 
fine  powder,  mixed  with  more  or  less  clay,  and  then  pressed  through 
perforated  plates,  from  which  the  "lead"  issues  in  tiny  rods.  These 
are  dried,  cut  into  the  proper  lengths,  baked  to  remove  all  traces  of 
moisture,  and  then  inserted  in  the  wooden  case. 

In  the  United  States  in  1902  over  four  million  pounds  of  graphite 
were  mined,  and  over  thirty-two  million  pounds  were  imported. 

Molten  iron  and  other  metals  dissolve  carbon,  and  when  the  metals 
cool  the  carbon  crystallizes  as  graphite.  Moissan  incidentally  obtained 
considerable  graphite  in  making  diamonds.  Artificial  graphite  is  now 
a  commercial  article  (see  Chapter  X). 

Amorphous  Carbon  is  a  broad  term,  including  all  vari- 
eties of  coal  and  charcoal,  lampblack,  and  gas  carbon. 
They  are  the  non-crystalline  forms  of  impure  carbon. 
The  word  amorphous  means  literally  "without  form,"  and 
it  is  often  used  to  designate  soft,  powdery,  and  uncrys- 
tallized  substances. 

Coal  is  a  term  applied  to  several  varieties  of  impure  carbon.  It  may 
be  regarded  as  the  final  product  derived  from  vegetable  matter  by  heat 
and  pressure  to  which  it  was  subjected  through  long  geological  periods. 

Ages  ago  the  vegetation  was  exceedingly  dense  and  luxuriant  upon 
land  slightly  raised  above  the  sea.  In  process  of  time  this  vegeta- 


FiG.  26.  —  Section  of  part  of  the  earth's  crust  near  Mauch  Chunk,  Penn., 
showing  layers  of  coal. 

tion  decayed,  accumulated,  and  slowly  became  covered  with  sand,  mud, 
and  water.  The  heat  of  the  earth  and  the  enormous  pressure  of  the 
overlaying  deposits  changed  the  vegetable  matter  into  more  or  .less 


Carbon  and  its  Oxides.  185 

impure  carbon.  This  series  of  geological  and  chemical  changes  was 
repeated,  and  as  a  result  we  find  in  the  earth  layers  or  seams  of  carbo- 
naceous matter  varying  in  thickness  and  composition  (Fig.  26).  These 
are  the  coal  beds. 

Coal  .beds  contain  proofs  of  their  vegetable  origin,  viz.,  impressions 
of  vines,  stems,  and  leaves  of  plants,  and  similar  vegetable  substances 


FIG.  27.  —  Fossil  found  in  a  FIG.  28.  —  Section  of  coal  as  seen  through 

coal  bed.  a  microscope. 

(Fig.  27).     A  thin  section  of  coal  examined  through  a  microscope  re- 
veals a  distinct  vegetable  structure  (Fig.  28). 

There  are  three  principal  kinds  of  coal,  (i)  Bitumi- 
nous or  soft  coal  is  used  to  make  illuminating  gas,  coke, 
and  as  a  fuel  for  steam ;  it  burns  with  a  smoky  flame,  and 
in  burning  produces  much  volatile  matter.  (2)  Anthra- 
cite coal  is  hard  and  lustrous.  It  ignites  with  difficulty, 
burns  with  little  or  no  flame,  and  produces  an  intense  heat. 
It  is  used  mainly  for  domestic  purposes, — heating  and 
cooking,  —  especially  in  eastern  United  States.  (3)  Lig- 
nite or  brown  coal  is  the  least  valuable  as  fuel.  It  often 
shows  the  woody  fiber  and  was  probably  formed  much 
later  than  the  other  varieties.  Peat,  strictly  speaking,  is 
not  coal,  though  it  is  used  as  fuel  in  some  places,  espe- 
cially in  Ireland  and  Holland.  It  is  formed  by  the  slow 


i86 


Descriptive  Chemistry. 


decay  of  roots  and  other  vegetable  matter  under  water, 
and  represents  an  early  stage  of  coal  formation. 

The  average  composition  of  different  kinds  of  coal  is 
seen  by  the  following  table  :  — 


,     KIND. 

CARBON. 

VOLATILE 
MATTER. 

ASH. 

WATER. 

Lignite    .... 

TO  Q 

2O  Q 

jO  2 

18 

->w-y 

Bituminous  .... 

74-53 

I5-I3 

10.34 

— 

Anthracite   .... 

91.64 

6.89 

1.47 

— 

Some  anthracite  coals  contain  as  much  as  95  to  99  per 
cent  of  carbon,  and  some  bituminous  coals  as  little  as  65 
per  cent.  Peat  and  wood  contain  still  less  carbon,  but 


FIG.  29.  — Coal  fields  in  the  United  States. 

more  volatile  matter.  The  volatile  matter  includes  nitro- 
gen, hydrogen,  and  sulphur.  These  facts  show  that  vege- 
table matter,  in  passing  through  the  changes  which  finally 


Carbon  and  its  Oxides. 


end  in  coal,  loses  volatile  matter,  Anthracite  coal,  which 
is  found  at  different  depths  and  associated  with  rocks  of 
different  ages,  shows  that  it  was  formed  from  the  bitumi- 
nous variety  by  the  great  pressure  caused  by  mountain 
building.  Hence  it  loses  volatile  matter  and  becomes  hard. 

Coal  is  widely  distributed  in  the  crust  of  the  earth,  but  the  deposits 
vary  in  extent  and  quality.  It  underlies  about  one  sixth  of  the  area  of 
the  United  States,  the  anthracite  variety  covering  less  than  five  hundred 
square  miles  in  eastern  Pennsylvania  (Fig.  29).  The  United  States 
now  leads  the  world  in  coal  production, 
furnishing  about  one  third  of  the  total 
supply.  England  for  many  years  headed 
the  list,  and  even  now  furnishes  a  large 
amount,  for  its  deposits  are  extensive 
(Fig.  30). 

Charcoal  is  a  variety  of  amor- 
phous carbon  obtained  by  heating 
wood,  bones,  ivory,  and  other 
organic  matter  in  closed  vessels, 
or  by  partially  burning  them  in 
the  air.  Th?  process  consists 
essentially  in  driving  off  the  vola- 
tile matter  and  retaining  the 
carbon. 

Wood  Charcoal  is  a  black, 
brittle  solid,  and  often  has  the 
form  of  the  wood  from  which  it 
is  made.  It  is  insoluble,  though 
its  mineral  impurities  may  be  removed  by  acids.  It 
burns  without  lame  or  much  smoke,  and  leaves  a  white 
ash.  •  The  compact  varieties  conduct  heat  and  electricity, 
but  porous  charcoal  is  a  poor  conductor.  It  resists  the 
action  of  many  chemicals;  hence  fence  posts,  telegraph 
poles,  and  wooden  piles  are  often  charred  before  being 


JKITISH 
COALFIELDS 


FlG.  30.  —  Coal  deposits  in  the 
British  Isles. 


1 88  Descriptive  Chemistry. 

put  into  the  ground.  Most  varieties  are  very  porous,  and 
when  thrown  upon  water  charcoal  floats,  owing  to  the 
presence  of  air  in  its  pores.  Its  porosity  makes  charcoal 
an  excellent  absorber  of  gases,  some  varieties  absorbing 
ninety  times  their  bulk  of  ammonia  gas.  Sewers  and  foul 
places  are  sometimes  purified  by  charcoal.  It  will  also 
absorb  colored  substances  from  solutions.  This  is  espe- 
cially true  of  animal  charcoal  (see  below).  Foul  air  and 
water  may  be  partially  purified  by  charcoal,  which  forms 
the  essential  part  of  many  water  filters  in  houses.  Char- 
coal used  for  such  a  purpose,  however,  must  be  renewed 
or  often  heated  to  redness;  otherwise  it  becomes  clogged 
and  contaminated.  Charcoal  is  never  pure  carbon,  the 
degree  of  purity  depending  upon  the  kind  of  wood  used, 
as  well  as  the  temperature  and  method  employed. 

Besides  the  uses  of  charcoal  mentioned  above,  it  is  used 
as  a  fuel,  in  the  manufacture  of  steel  and  of  gunpowder, 
and  as  a  medicine.  It  reduces  oxides  when  heated  with 
them,  thus  — 

2  CuO      +     C      =  2  Cu  +          CO2 
Copper  Oxide      Carbon      Copper      Carbon  Dioxide 

Wood  charcoal  is  made  either  in  a  charcoal  pit  or  kiln,  or  in  a  large 
retort.  Where  wood  is  plentiful,  it  is  loosely  piled  into  the  shape 
shown  in  Figure  31,  and  covered  with  turf  to  prevent  free  access  of  air, 
though  small  holes  are  left  at  the  bottom  and  a  larger  one  at  the  top  of 
a  central  flue,  so  that  sufficient  air  can  pass  through  the  pile.  The 
wood  is  lighted,  and  as  it  slowly  burns  care  is  taken  to  regulate  the 
supply  of  air,  so  that  the  wood  will  smolder  but  not  burn  up.  The 
volatile  matter  escapes  and  charcoal  remains,  the  average  yield  being 
about  20  per  cent  of  the  weight  of  the  wood.  This  method  is  crude, 
uncertain,  and  wasteful.  Much  charcoal  is  now  made  by  heating 
wood  in  closed  retorts,  no  air  whatever  being  admitted.  By  this 
method,  which  is  called  dry  or  destructive  distillation,  the  yield  of 
charcoal  is  30  per  cent  and  all  the  volatile  matter  is  saved.  In  the 


Carbon  and  its  Oxides.  189 

ordinary  combustion  of  wood,  the  hydrogen  forms  water  and  the  oxy- 
gen forms  carbon  dioxide ;  but  in  dry  distillation,  where  no  oxygen  is 
present,  much  of  the  hydrogen  forms  volatile  compounds  with  the  car- 
bon and  oxygen.  Among  these  volatile  products  are  methyl  alcohol 


FIG.  31.  —  Wood  arranged  for  burning  into  charcoal. 

and  acetic  acid.  These  are  commercial  substances,  and  contribute  to 
the  profit  of  the  process.  More  or  less  charcoal  is  obtained  by  heating 
any  compound  of  carbon,  e.g.  sugar  or  starch,  the  charring  being  a  test 
for  carbon. 

Animal  Charcoal  or  Bone  Black  is  made  by  heating  bones  in  a  closed 
vessel,  and  by  heating  a  mixture  of  blood  and  sodium  carbonate.  It 
contains  only  about  10  per  cent  of  carbon,  but  this  carbon  is  dis- 
tributed throughout  the  porous  mineral  matter  of  the  bone,  which  is 
almost  entirely  calcium  phosphate.  Under  the  name  of  ivory  black, 
animal  charcoal  is  used  as  a  pigment,  especially  in  making  shoe-black- 
ing. It  is  extensively  used  to  remove  the  color  from  sugar  sirups,  oils, 
and  other  liquids  colored  by  organic  matter. 

Coke  is  made  by  expelling  the  volatile  matter  from  soft 
coal,  somewhat  as  charcoal  is  made  from  wood.  It  is  left 
in  the  retorts  when  coal  is  distilled  in  the  manufacture  of 
illuminating  gas.  On  a  large  scale  it  is  made  by  heating 
a  special  grade  of  soft  coal  in  huge  brick  ovens,  shaped 
like  a  beehive,  from  which  air  is  excluded  after  combus- 
tion begins.  Sometimes  the  coke  is  made  in  closed  retorts 
constructed  so  as  to  save  the  by-products,  —  ammonia,  tar, 


190  Descriptive  Chemistry. 

organic  compounds,  and  combustible  gases.  This  method 
not  only  yields  more  coke,  but  is  also  more  profitable  be- 
cause the  by-products  are  sold  and  the  combustible  gas  is 
used  to  heat  the  retorts.  Coke  is  a  grayish,  porous  solid, 
harder  and  heavier  than  charcoal.  It  burns  with  no  smoke 
and  a  feeble  flame.  It  contains  about  90  per  cent  of  car- 
bon, the  rest  being  the  mineral  matter  originally  in  the  coal. 

Immense  quantities  of  coke  are  used  in  the  manufacture  of  iron  and 
steel.  It  is  superior  to  coal  for  this  purpose,  because  it  gives  a  greater 
heat  when  burned,  reduce's  oxides  easily,  and  contains  little  or  no 
sulphur  or  other  substances  harmful  in  the  iron  industries.  Coke  is 
the  fuel  used  in  making  nine  tenths  of  the  pig  iron  in  the  United 
States,  and  over  twelve  million  tons  (or  about  three  fourths  of  the 
total  amount)  are  made  annually  in  the  Connellsville  district,  near 
Pittsburg,  Pennsylvania. 

Gas  Carbon  is  amorphous  carbon  which  is  gradually  deposited  upon 
the  inside  of  the  retorts  used  in  the  manufacture  of  illuminating  gas. 
It  is  a  black,  heavy,  hard  solid,  and  is  almost  pure  carbon.  It  is  a  good 
conductor  of  electricity,  and  is  extensively  used  for  the  manufacture  of 
the  carbon  rods  of  electric  lights  and  for  plates  of  electric  batteries. 

Lampblack  is  prepared  by  burning  oil  or  oily  substances  rich  in 
carbon  in  a  limited  supply  of  air.  The  dense  smoke,  which  is  mainly 
finely  divided  carbon,  is  passed  through  a  series  of  condensing  cham- 
bers, where  it  is  collected  upon  coarse  cloth  or  a  cold  surface.  Its 
formation  is  illustrated  on  a  small  scale  by  a  smoking  lamp,  and  the 
soot  deposited  is  the  same  as  lampblack.  Lampblack  is  one  of  the 
purest  forms  of  amorphous  carbon,  and  it  is  used  in  making  printer's 
ink  and  certain  black  paints. 

Allotropism.  —  Diamond,  graphite,  and  amorphous  car- 
bon, though  exhibiting  essentially  different  properties,  are 
identical  in  composition.  All  are  carbon.  They  can  be 
changed  into  one  another,  the  amorphous  form  into  graph- 
ite and  finally  into  diamond  and  the  diamond  into  amor- 
phous carbon.  Each  burns  in  oxygen  and  the  product  is 
carbon  dioxide.  Furthermore,  the  same  weight  of  each 


Carbon  and  its  Oxides.  191 

forms  the  same  weight  of  carbon  dioxide,  i.e.  when  12 
gm.  of  each  are  burned,  44  gm.  of  carbon  dioxide  are 
always  produced.  There  is  no  doubt  about  their  identity, 
though  no  one  has  explained  it.  The  property  of  assum- 
ing more  than  one  elementary  form  is  called  allotropism 
or  allotropy  (from  Greek  words  meaning  another  form). 
The  more  uncommon  form  is  called  an  allotrope  or  an 
allotropic  modification  of  the  other.  It  is  believed  by  some 
that  allotropism  is  due  to  a  difference  in  the  number  of 
atoms  in  a  molecule  of  the  element. 

OXIDES    OF    CARBON. 

Carbon  and  Oxygen  do  not  unite  at  the  ordinary  tem- 
perature. But  when  carbon  is  heated  in  air,  in  oxygen,  or 
with  some  oxides,  carbon  dioxide  (CO2)  is  formed ;  if  the 
supply  of  oxygen  is  limited,  then  carbon  monoxide  (CO) 
is  formed. 

Occurrence  and  Formation  of  Carbon  Dioxide.  — The 

occurrence  of  carbon  dioxide  in  the  atmosphere  and  in 
many  natural  waters  has  already  been  mentioned.  It  is 
the  main  product  of  ordinary  combustion,  respiration  of 
animals,  and  decay.  In  all  these  processes  the  carbon 
comes  from  organic  matter,  while  the  oxygen  comes  from 
the  air,  from  the  organic  matter,  or  from  both. 

Ordinary  combustion  is  a  chemical  combining  of  carbon 
and  oxygen.  Hence,  when  carbon  or  a  substance  contain^ 
ing  it  is  burned,  carbon  dioxide  is  formed.  The  equation 
for  this  change  is  — 

C         +         02  C02 

Carbon  Oxygen  Carbon  Dioxide 

Carbon  dioxide  is  formed  by  the  combustion  of  such  com- 
mon substances  as  wood,  coal,  charcoal,  coke,  oils,  waxes, 


192  Descriptive  Chemistry. 

cotton,  bone,  starch,  sugar,  meat,  bread,  alcohol,  camphor, 
and  illuminating  gas. 

The  continuous  oxidation  of  the  tissues  and  foods  in  the 
body  produces  carbon  dioxide  (see  Relation  of  Oxygen  to 
Life).  And  if  we  exhale  the  breath  through  a  glass  tube 
into  limewater,  the  carbon  dioxide  which  is  in  the  breath 
turns  the  limewater  milky  —  the  usual  test  for  carbon 
dioxide.  The  equation  for  the  change  is  — 

CO2  +     Ca(OH)2    =          CaCO3          +     H2O 

Carbon  Dioxide  Limewater  Calcium  Carbonate 

When  vegetable  and  animal  matter  decays,  carbon 
dioxide  is  formed.  Many  kinds  of  organic  matter  fer- 
ment, especially  those  containing  sugar.  By  alcoholic 
fermentation  the  sugar  changes  into  carbon  dioxide  and 
alcohol  (see  Alcohol),  thus  — 

C6H12°6        =  2CO2  +  2C2H6O 

Sugar  Carbon  Dioxide  Alcohol 

The  Preparation  of  Carbon  Dioxide  is  usually  accom- 
plished by  the  interaction  of  a  carbonate  and  an  acid. 
Calcium  carbonate  (limestone  or  marble)  and  hydrochloric 
acid  are  usually  used.  The  operation  may  be  easily  per- 
formed in  any  glass  vessel  by  pouring  the  acid  upon  the 
carbonate.  The  equation  for  the  chemical  change  is  — 

CaCOg     +     2HC1     =     CO2     +     CaCl2     +      H2O 
Calcium  Carbon  Calcium 

Carbonate  Dioxide  Chloride 

This  gas  may  also  be  prepared  by  heating  matter  con- 
taining carbon,  or  by  strongly  heating  carbonates  (as  in 
making  lime),  thus  — 

CaCOg  CO2         +     CaO 

Calcium  Carbonate        Carbon  Dioxide          Lime 


Carbon  and  its  Oxides. 

Properties  of  Carbon  Dioxide.  —  This  gas  has  many 
important  properties  besides  those  mentioned  under  The 
Atmosphere.  It  has  a  slight  taste  and  odor,  but  no  color. 
It  is  one  and  a  half  times  heavier  than  air,  and  a  liter 
under  standard  conditions  weighs  1.977  gm.  On  ac- 
count of  its  weight  it  can  be  collected  by  downward  dis- 
placement and  poured  from  one  vessel  to  another.  For 
the  same  reason,  it  is  often  found  at  the  bottom  of  old  or 
deep  wells,  in  some  valleys  near  lime  kilns  or  volcanoes, 
and  in  mines  after  explosions.  At  the  ordinary  tempera- 
ture and  pressure,  water  dissolves  its  own  volume  of 
carbon  dioxide.  Under  increased  pressure  more  gas  dis- 
solves, which  escapes  readily  when  the  pressure  is  re- 
moved. Hence  "  soda  water,"  which  is  made  by  forcing 
carbon  dioxide  into  water,  effervesces  and  froths  when 
drawn  from  the  soda  fountain.  Many  natural  waters  and 
manufactured  beverages  (such  as  champagne  and  beer) 
sparkle  and  effervesce  for  the  same  reason.  This  gas  may 
be  liquefied  by  subjecting  it  to  high  pressure  and  low 
temperature.  It  was  first  liquefied  by  Faraday  by  the 
method  used  for  chlorine.  Liquid  carbon  dioxide  is  now 
made  in  large  quantities  by  forcing  the  gas  into  steel 
cylinders  by  powerful  pumps,  the  gas  being  obtained  in 
many  cases  from  the  fermenting  vats  of  breweries. 
When  a  cylinder  of  liquid  .carbon  dioxide  is  opened,  the 
liquid  evaporates  so  rapidly  that  a  portion  of  it  becomes 
a  white,  snowlike  solid.  Both  the  liquid  and  solid  carbon 
dioxide  are  articles  of  commerce,  and  are  sometimes 
used  to  prepare  "soda  water,"  to  extinguish  fires,  to 
improve  wines,  and  to  produce  very  low  temperatures. 
Carbon  dioxide  extinguishes  burning  objects,  such  as  a 
blazing  stick  or  lighted  candle;  indeed,  air  containing 
from  2.5  to  4  per  cent  of  carbon  dioxide  will  extinguish 


194  Descriptive  Chemistry. 

small  flames.  Hence  the  gas  is  often  used  to  extinguish 
fires.  Many  small  fire  extinguishers  contain  sodium 
carbonate  and  sulphuric  acid,  so  arranged  that  when 
desired,  carbon  dioxide  gas  may  be  generated  from  them 
under  pressure.  A  stream  of  the  gas  forced  upon  a 
small  blaze  will  often  prevent  a  serious  fire.  In  other 
forms,  the  carbon  dioxide,  which  is  similarly  generated, 
forces  water  from  the  extinguisher. 

Relation  of  Carbon  Dioxide  to  Life.  — Animals  die  when 
put  into  carbon  dioxide.  It  cuts  off  the  supply  of  oxygen 
as  water  does  from  a  drowning  man.  The  presence  of  a 
small  quantity  in  the  air  is  objectionable,  since  it  is  said  to 
produce  headache  and  drowsiness;  but  much  of  the  dis- 
comfort felt  in  badly  ventilated  rooms  and  attributed  to 
carbon  dioxide  is  doubtless  due  to  water  vapor,  and  to 
poisonous  substances  produced  from  the  organic  mat- 
ter exhaled  from  the  lungs.  On  the  other  hand,  carbon 
dioxide  is  an  essential  food  of  plants.  Through  their 
leaves  and  other  green  parts  they  absorb  carbon  dioxide 
from  the  atmosphere,  decompose  it,  reject  the  oxygen,  and 
store  up  the  carbon  in  the  form  of  starch.  The  sunlight 
and  the  green  coloring  matter  aid  the  plant  in  manufac- 
turing its  food  out  of  the  water  (obtained  through  the  roots 
from  the  soil)  and  the  carbon  of  the  carbon  dioxide  ob- 
tained from  air.  Plants  thus  serve  to  keep  the  atmosphere 
free  from  an  excess  of  carbon  dioxide,  the  proportion 
present  in  the  air  being  very  small  and  practically  con- 
stant. 

Carbonic  Acid.  —  Carbon  dioxide  gas  is  often  called  carbonic  acid  gas, 
or  simply  carbonic  acid.  It  is  believed  that  carbon  dioxide,  when  passed 
into  water,  combines  with  the  water  and  forms  a  weak,  unstable  acid, 
which  is,  strictly  speaking,  carbonic  acid.  The  equation  for  this 
change  is  — 


Carbon  and  its  Oxides.  195 

CO2    +    H20    =    H2CO3 
Carbon  Dioxide  Carbonic  Acid 

Such  a  solution  reddens  blue  litmus  and  decolorizes  pink  phenolphthal- 
ein.  Carbonic  acid  has  never  been  obtained  free,  and  is  so  unstable 
that  it  easily  breaks  up  by  gentle  heat  into  carbon  dioxide  and  water, 
thus  — 

H,CO3  =  CO2  +  H20. 

Carbon  dioxide  is  sometimes  called  carbonic  anhydride,  to  denote  its 
relation  to  the  acid. 

Carbonates  are  salts  corresponding  to  the  unstable 
carbonic  acid.  They  are  stable  compounds.  The  most 
abundant  natural  carbonates  are  those  of  calcium,  magne- 
sium, and  iron.  Immense  quantities  of  sodium  and  potas- 
sium carbonates  are  manufactured. 

A  few  carbonates  are  formed  by  direct  combination  of  an  oxide  and 
carbon  dioxide,  but  most  of  them  are  formed  by  passing  carbon  dioxide 
into  the  corresponding  hydroxide,  thus  — 

CO2     +  Ca(OH)2  CaCO3  +     H2O 

Calcium  Hydroxide     Calcium  Carbonate 

Many  carbonates  are  insoluble  in  water,  e.g.  calcium  carbonate,  the 
test  for  carbon  dioxide  depending  upon  this  fact.  Others,  e.g.  sodium 
and  potassium  carbonate,  are  very  soluble.  There  are  two  classes  of 
carbonates,  the  normal  and  the  acid.  Normal  sodium  carbonate  is 
Na2CO3,  and  acid  sodium  carbonate  is  HNaCO3.  The  latter  is  often 
called  sodium  bicarbonate.  Normal  calcium  carbonate  is  CaCO3,  and 
acid  calcium  carbonate  is  H2Ca(CO3)2 ;  4he  latter  is  unstable,  and  is 
easily  decomposed  by  heat  into  normal  calcium  carbonate. 

Composition  of  Carbon  Dioxide.  — If  a  known  weight  of  pure  car- 
bon, such  as  diamond  or  graphite,  is  burned  in  oxygen,  it  is  found  that 
for  12  parts  of  carbon  used  there  are  44  parts  of  carbon  dioxide  formed. 
Hence  12  parts  of  carbon  unite  with  32  parts  of  oxygen.  The  vapor 
density  of  the  gas  is  22,  and  the  molecular  weight  must  be  44.  These 
facts  necessitate  the  formula  CO2. 


196  Descriptive  Chemistry. 

History  of  Carbon  Dioxide. — This  gas  was  described  in  the  seven- 
teenth century  by  Van  Helmont,  who  called  it  gas  sylvestre.  He 
prepared  it  by  the  interaction  of  acids  and  carbonates,  detected  it  in 
mineral  water,  and  observed  its  formation  during  combustion  and  fer- 
mentation, as  well  as  its  action  on  animals  and  flames.  Black,  in  1755, 
showed  that  carbon  dioxide  is  essentially  different  from  ordinary  air  and 
that  the  gas  is  readily  obtained  from  magnesium  and  calcium  carbonates. 
Since  the  gas  was  combined  or  "  fixed  "  in  these  substances,  he  called 
the  gas  fixed  air.  His  work  was  verified  in  1774  by  Bergman,  who 
called  the  gas  acid  of  air.  Lavoisier  first  proved  it  to  be  an  oxide  of 
carbon. 

Carbon  Monoxide  is  formed  when  carbon  is  burned  in  a 
limited  supply  of  air,  thus  — 

C         +         O  CO 

Carbon  Oxygen          Carbon  Monoxide 

If  carbon  dioxide  is  passed  over  heated  charcoal,  the  prod- 
uct is  carbon  monoxide.  That  is,  carbon  reduces  carbon 
dioxide  to  carbon  monoxide,  the  equation  for  the  change 
being  •— 

C02         +         C  2  CO 

Carbon  Monoxide 

This  chemical  change  takes  place  in  every  coal  fire.  The 
oxygen  of  the  air  entering  the  bottom  of  the  fire  unites  with 
the  carbon  to  form  carbon  dioxide  ;  the  latter  gas  in  passing 
through  the  hot  carbon  of  the  fire  is  reduced  to  carbon 
monoxide.  Some  of  the  carbon  monoxide  escapes  and 
some  burns  with  a  flickering  bluish  flame  on  the  top  of 
the  fire. 

If  steam  is  passed  over  red-hot  coke  or  charcoal,  a  mixture  of  carbon 
monoxide  and  hydrogen  is  produced.  This  mixture  enriched  by  vapor 
from  oils  is  known  as  water  gas  (see  Water  Gas) . 

Carbon  monoxide  is  usually  prepared  by  gently  heating 
a  mixture  of  oxalic  acid  and  sulphuric  acid  in  a  flask,  and 


Carbon  and  its   Oxides.  197 

collecting  the  gaseous  product  over  water.  The  oxalic  acid 
decomposes  thus  — 

C2H2O4     =  CO  +         CO2         +      H2O 

Oxalic  Acid  Carbon  Monoxide          Carbon  Dioxide 

The  carbon  dioxide  may  be  removed  by  passing  the  mixed 
gases  through  a  solution  of  sodium  hydroxide. 

Carbon  monoxide  is  a  gas  without  color,  odor,  or  taste,  and 
is  only  slightly  soluble  in  water.  It  burns  with  a  bluish 
flame,  forming  carbon  dioxide,  thus  — 

2  CO  +          .O2  2CO2 

Carbon  Monoxide  Carbon  Dioxide 

Carbon  monoxide  is  extremely  poisonous,  and  it  is  doubly 
dangerous  because  its  lack  of  odor  prevents  its  detection  in 
time  to  escape  its  stupefying  effect.  Many  deaths  have 
been  caused  by  breathing  air  containing  it.  Carbon  mo- 
noxide forms  a  compound  with  one  of  the  constituents  of 
the  blood,  and  those  who  have  been  poisoned  by  it  cannot 
be  revived  by  air,  as  in  the  case  of  suffocation  by  carbon 
dioxide.  It  is  a  constituent  of  ordinary  illuminating  gas, 
and  care  should  always  be  taken  to  prevent  the  escape  of 
illuminating  gas  (as  well  as  the  gas  from  a  coal  stove  or 
furnace)  into  rooms  occupied  by  human  beings.  At  a  high 
temperature  carbon  monoxide  unites  easily  with  oxygen, 
and  is,  therefore,  an  important  agent  in  the  reduction  of 
iron  ores  in  the  blast  furnace.  This  action  might  be  rep- 
resented thus  — 

Fe203     +  3  CO          =      2Fe      +        3  CO2 

Iron  Oxide  Carbon  Monoxide  Iron  Carbon  Dioxide 

Carbon  monoxide,  which  is  sometimes  called  carbonic  oxide,  forms  no 
acid  and  therefore  no  salts.  It  does  not  make  limewater  milky,  thus 
being  readily  distinguished  from  carbon  dioxide.  Its  blue  flame  dis- 


198  Descriptive  Chemistry. 

tinguishes  it  from  all  other  gases  which  burn.  It  unites  directly  with 
chlorine  to  form  carbonyl  chloride  (phosgene,  COC12),  and  with  some 
metals,  forming  metallic  carbonyls,  e.g.  nickel  carbonyl  (Ni(CO)4). 

Cyanogen  is  a  compound  of  carbon  and  nitrogen  having 
the  composition  corresponding  to  the  formula  (CN)2.  It 
is  a  colorless  gas,  has  the  odor  of  peach  kernels,  is  exceed- 
ingly poisonous,  and  burns  with  a  purplish  flame.  It  may 
be  prepared  by  heating  mercuric  cyanide  (Hg(CN)2). 
Cyanogen  is  a  radical,  and  in  compounds  it  acts  like  an 
element.  Its  corresponding  acid  is  hydrocyanic  or  prus- 
sic  acid  (HCN).  This  acid  is  prepared  by  heating  a 
cyanide  with  sulphuric  acid,  "just  as  hydrochloric  acid  is 
obtained  from  a  chloride.  The  solution  smells  like  peach 
kernels,  and  is  one  of  the  most  deadly  of  all  known  poisons. 
Potassium  cyanide  is  a  white,  deliquescent  solid.  It  is 
a  deadly  poison.  Large  quantities  are  used  in  gold  and 
silver  plating  and  in  the  "  cyanide  process  "  of  extracting 
gold  from  its  ores,  as  described  under  that  metal.  Other 
cyanogen  compounds  are  cyanic  acid  (CNOH),  sulpho- 
cyanic  acid  (CNSH),  and  potassium  sulphocyanate 
(CNSK).  The  last  is  a  white,  crystallized  salt,  which 
produces  a  beautiful  red  solution  when  added  to  certain 
soluble  iron  compounds,  and  is  therefore  used  to  detect 
this  metal.  Salts  of  complex  acids  related  to  hydrocyanic 
acid  are  used  in  dyeing,  many  being  prepared  from  the 
most  common  one  —  potassium  ferrocyanide  or  yellow 
prussiate  of  potash.  They  will  be  described  in  the  chap- 
ter on  Iron. 

EXERCISES. 

1.  What  is  the  symbol  and  atomic  weight  of  carbon? 

2.  In  what  forms  does  free  carbon  occur  in  nature?     Name  ten  famil- 
iar solids,  three  liquids,  and  two  gases  containing  carbon.     What  pro- 
portion of  the  earth's  crust  is  carbon? 


Carbon  and  its  Oxides.  199 

3.  What  is  diamond?  How  could  the  correctness  of  your  answer  be 
shown?     State  (a}  the  source,  (b)  the  properties,  and  (c)  the  uses  of 
diamonds.     Give  a  brief  account  of  one  or  more  famous  diamonds. 

4.  What  is  graphite?     What  is  its  chemical  relation  to  diamond, 
and  how  could  this  relation  be  proved  ?     State  (a)  the  source,  (b}  the 
properties,  and  (c)  the  uses  of  native  graphite. 

5.  What  is  {a}  black  lead,  (£)  plumbago,  (c)  bort,  (d)  carbonado, 
(e)  native  graphite,  (/)  artificial  graphite? 

6.  Give  a  brief  account  of  the  manufacture  of  lead  pencils.     What 
is  the  literal  meaning  of  graphite? 

7.  Review  artificial  graphite  (see  Chapter  X). 

8.  What  does  the  term  amorphous  carbon  include?     Does  the  car- 
bon in  these  impure  forms  differ  chemically  from  diamond  and  graphite? 

9.  How  was  coal  formed?     Give  several  proofs  of  its  origin.     State 
the  properties  and  uses  of  (a)  bituminous  coal,  (b}  anthracite  coal,  and 
(V)  lignite.      What  besides  carbon  does  it  contain?     Where  is  coal 
found  ? 

10.  WThat  is  charcoal  ?  State  (a}  the  properties,  and  (b}  the  uses  of 
wood  charcoal.  Give  a  brief  account  of  both  methods  of  preparing  wood 
charcoal.  State  the  preparation,  properties,  and  uses  of  animal  charcoal. 

u.  What  is  coke?  How  is  it  made?  What  are  its  properties? 
How  is  it  related  to  the  iron  industries? 

12.  What  is  gas  carbon?     What  is  its  source?     State  its  properties 
and  uses. 

13.  What  is  lampblack?  State  its  method  of  preparation,  properties, 
and  uses. 

14.  Define  and  illustrate  (a}  amorphous,  and  (b}  allotropism. 

15.  Develop  the  topics:  (a}  carbon  is  a  reducing  agent,  (b}  carbon 
monoxide  is  a  reducing  agent,  (c)  diamond,  graphite,  and  pure  amor- 
phous carbon  illustrate  allotropism. 

1 6.  What  is  (a)  hard  coal,  (b)  soft  coal,  (c)  peat,  (d}  boneblack, 
(e)  soot,  (/)  lampblack,  (g)  lignite,  (h)  electric  light  carbon? 

17.  Give  the  names  and  formulas  of  the  two  oxides  of  carbon.     How 
is  each  formed  from  carbon  and  oxygen? 

1 8.  Describe  the  occurrence  and  formation  of  carbon  dioxide.  What 
is  always  obtained  by  burning  a  substance  containing  carbon  ?     Give 
the  simplest  equation  for  this  chemical  change. 

19.  Describe  fully  the  action  of  carbon  dioxide  on  limewater.     Give 
the  equation  for  the  reaction. 


2OO  Descriptive  Chemistry. 

20.  What  is  the  relation  of  carbon  dioxide  to  (a}  respiration,  (b)  fer- 
mentation of  sugar,  (c)  decay,  (d)  making  lime  ? 

21.  What  is  the  test  for  (a)  carbon,  (£)  carbon  monoxide,  (c}  car- 
bon dioxide? 

22.  Describe  the  usual  method  of  preparing  carbon  dioxide.     Give 
the  equation  for  the  reaction.     State  its  properties. 

23.  Describe  liquid  and  solid  carbon  dioxide.     How  are  they  pre- 
pared ?     For  what  are  they  used  ? 

24.  What  is  the  relation  of  carbon  dioxide  to  animal  and  to  plant 
life? 

25.  State  fully  the  relation  of  carbon  dioxide  to  the  unstable  acid 
H.,CO3.     Give  the  equations  for  the  formation  and  decomposition  of 
this  acid. 

26.  What  are  carbonates?     Name  three.     How  are  they  formed? 
What  are  their  properties  ? 

27.  What  is  (a)  "  soda  water,"  (£)  carbonated  water,  (c)  carbonic 
acid,  (d)  carbonic  oxide,   (e)  carbonic  anhydride,  (/)  limestone  or 
marble? 

28.  What  is  the  difference  between  (a}  sodium  carbonate  and  sodium 
bicarbonate,  and  (b)  calcium  carbonate  and  acid  calcium  carbonate  ? 

29.  Why  is  (a)  CO2  the  formula  of  carbon  dioxide,  and  (b}  CO  of 
carbon  monoxide? 

30.  State  briefly  the  history  of  carbon  dioxide. 

31.  Give   a   brief  account  of   (a}   Black,   (b)  Van   Helmont,  and 
(c)  Bergman. 

32.  Illustrate   the  law  of  multiple   proportions   by  the   oxides  of 
carbon. 

33.  Give  the  equations  for  (a}  the  oxidation  of  carbon  to  carbon 
monoxide,   (^)  the  reduction  of  carbon  dioxide  to  carbon*  rnonoA^-.. 

34.  How  is  carbon  monoxide  (#)  formed,  and  (<£)  usually  prepared  ? 

35.  What  is  the  relation  of  carbon  monoxide  to  water  gas? 

36.  What  are  the  properties  of  carbon  monoxide? 

37.  Illustrate  Gay-Lussac's  law  by  the  combustion  of  carbon  mo- 
noxide (2  CO  +  O.,  =  2  CO2) . 

38.  Illuminating  gas,  water  gas,  and  the  gas  which  escapes  from  a 
coal  fire  are  poisonous.     Why? 

39.  What  is  cyanogen?     Hydrocyanic   acid?     Describe  potassium 
cyanide.     For  what  is  it  used?     Describe  ootassium  sulphocyanate. 
State  its  chief  use. 


Carbon  and  its  Oxides.  101 

40.  The  specific  gravity  of  charcoal  is  about  1.5.     Why  does  it  float 
on  water? 

41.  How  can  carbon  monoxide  and  carbon  dioxide  be  changed  into 
each  other? 

42.  Review  (a)  combustion,  (£)  solution  of  gases  (especially  carbon 
dioxide)  in  water,  (c}  respiration. 

43.  State  and  explain  the  various  chemical  changes  which  occur  from 
the  entrance  of  oxygen  (in  the  air)  below  the  grate  of  a  red-hot  coal 
fire  to  the  end  of  the  burning  of  the  carbon  monoxide  at  the  top  of  the 
coal. 

PROBLEMS. 

1 .  How  many  grams  of  calcium  carbonate  are  needed  to  prepare 
132  gm.  of  carbon  dioxide  ? 

2.  What  weight  of  carbon  burned  in  air  will  produce  n  gm.  of 
carbon  dioxide  ? 

3.  Calculate  the  percentage  composition  of  (a}  calcium  carbonate, 
(£)  carbon  monoxide,  (c)  carbon  dioxide,  (d)  magnesium  carbonate. 

4.  What  per  cent  of  carbon  (by  weight)  is  contained  in  carbon 
monoxide  and  in  carbon  dioxide  ? 

5.  If  20  gm.  of  carbon  are  heated  in  the  presence  of  44  gm.  of 
carbon  dioxide,  (a)  what  weight  of  carbon  monoxide  is  formed,  and  (<£) 
what  weight,  if  any,  of  carbon  remains  ? 

6.  How  many  liters  of  carbon  dioxide  must  be  passed  over  red-hot 
charcoal  to  yield  84  gm.  of  carbon  monoxide  ? 

7.  How  much  carbon  dioxide  («)  by  weight  and  (£)  by  volume  is 
in  the  air  of  a  room  6  m.  long,  4  m.  wide,  and  3  m.  high,  if  there  is 
i  vol.  of  carbon  dioxide  in  1000  vol.  of  air  ? 

8.  What  weight  of  water  must  be  decomposed  to  furnish  enough 
oxygen  to  form  (with  pure  carbon)  44  gm.  of  carbon  dioxide  ? 

9.  How  many  grams  of  calcium  carbonate  will  produce  15  1.  of 
carbon  dioxide  ? 

10.    If  a  piece  of  pure  graphite  weighing  7  gm.  is  burned  in  oxygen, 
what  volume  of  carbon  dioxide  is  formed  ? 


CHAPTER   XV. 

HYDROCARBONS  —  METHANE  —  ETHYLENE  —  ACETYLENE 
-ILLUMINATING  GAS  — FLAME  — BUNS  EN  BURNER  - 
OXIDIZING  AND  REDUCING  FLAMES. 

Hydrocarbons  are  compounds  of  carbon  and  hydrogen. 
They  number  about  two  hundred,  and  their  properties 
vary  between  wide  limits.  They  are  found  in  petroleum 
and  its  products  (kerosene,  naphtha,  lubricating  oils,  par- 
affin wax,  etc.),  in  coal  tar,  in  coal  gas  and  natural  gas, 
and  in  some  essential  oils,  such  as  turpentine.  On  a  large 
scale  they  are  prepared  by  the  destructive  distillation  of 
petroleum,  wood,  coal,  and  coal  tar.  Indirectly  the  hydro- 
carbons are  the  source  of  many  other  compounds  of  car- 
bon, which  are  extensively  used  in  numerous  industries. 

The  existence  of  so  many  hydrocarbons  is  due  to  the  fact  that  atoms 
of  carbon  have  power  to  unite  with  themselves.  This  property  gives 
rise  to  compounds  which  form  natural  groups  or  series.  Simple  rela- 
tions exist  between  many  hydrocarbons,  especially  between  members 
of  the  same  series.  The  consecutive  members  of  a  series  differ  in  com- 
position by  CH2.  Thus,  in  the  methane  series,  methane  is  CH4  and 
ethane  is  C2H6 ;  in  the  ethylene  series,  ethylene  is  C2H4  and  propylene 
is  C3H6;  in  the  acetylene  series,  acetylene  is  C2H2  and  allylene  is 
C3H4 ;  and  in  the  benzene  series,  benzene  is  C6HG  and  toluene  is  C-H8. 
These  series  are  called  homologous  series. 

Methane  is  found  in  coal  mines,  being  a  gaseous  prod- 
uct of  the  processes  which  changed  vegetable  matter 
into  coal.  It  is  called  fire  damp  by  miners.  It  is  also 
formed  in  marshy  places  by  the  decay  of  vegetable  matter 
under  water,  and  is  therefore  often  called  marsh  gas. 


Methane.  203 

It  is  a  constituent  of  natural  gas  and  petroleum,  and  forms 
a  large  proportion  of  the  illuminating  gas  obtained  by 
heating  coal. 

Methane  is  usually  prepared  in  the  laboratory  by  heating  a  mixture 
of  sodium  acetate,  sodium  hydroxide,  and  quicklime  in  a  hard  glass  or 
metal  vessel,  and  collecting  the  gaseous  product  over  water.  It  may 
also  be  prepared  by  the  interaction  of  aluminium  carbide  and  water, 
thus  — 

A13C4  +i2H20=      3CH4      +  4A1(OH)8 

Aluminium  Carbide         Water         Methane         Aluminium  Hydroxide 

Methane  has  no  color,  taste,  or  odor.  It  burns  with  a 
pale,  luminous  flame.  A  mixture  of  methane  with  oxygen 
or  air  explodes  violently  when  ignited  by  a  spark  or  flame. 
Terrible  disasters  occur  in  coal  mines  from  this  cause.  The 
products  of  the  explosion  are  carbon  dioxide  and  water, 
thus-  CR4  +  2Q2  =  co?  +  2H20 
Methane  Oxygen  Carbon  Dioxide  Water 

The  carbon  dioxide,  called  choke  damp  or  black  damp  by 
the  miners,  often  suffocates  those  who  escape  from  the 
explosion. 

Other  members  of  the  methane  series  are  ethane  (C2H6),  propane 
(C3H8),  butane  (C4H10).  This  series  is  also  called  the  paraffin  series, 
on  account  of  the  chemical  indifference  of  its  members.  It  has  the 
general  formula  CnH2n  +  2-  Butane  and  the  succeeding  fifteen  or 
twenty  members  are  liquids,  and  the  highest  members  are  solids. 

Chlorine  and  hydrocarbons  interact,  that  is,  chlorine  replaces  hydro- 
gen, atom  for  atom.  Thus  — 

CH4    +     2C1    =       CH3Cr        +    HC1 
Methane  Chlormethane 

This  chemical  change  is  called  substitution,  and  illustrates  one  of  the 
methods  used  in  preparing  derivatives  of  carbon  known  as  substitution 
products.  The  paraffins  are  saturated  hydrocarbons.  This  means 
that  the  carbon  in  them  is  saturated,  so  to  speak,  with  hydrogen,  and 
has  no  tendency  to  unite  directly  with  more  atoms  of  hydrogen  or 
other  elements. 


204  Descriptive  Chemistry. 

Ethylene  or  olefiant  gas  is  formed  by  the  destructive 
distillation  of  wood  and  coal.  It  is  usually  prepared  by 
heating  a  mixture  of  concentrated  sulphuric  acid  and  ethyl 
alcohol,  and  collecting  the  gas  over  water.  The  alcohol 
decomposes  into  ethylene  and  water,  the  latter  being  ab- 
sorbed by  the  sulphuric  acid.  The  essential  change  is 
represented  thus  — 

C^WgO      =      C^H^     -f-      H^O 
Alcohol  Ethylene 

Ethylene  is  a  colorless  gas,  and  has  a  pleasant  odor.  It 
can  be  condensed  to  a  liquid,  which  by  evaporation  pro- 
duces a  temperature  as  low  as  —  140°  C.  It  burns  with  a 
bright,  yellow  flame,  and  is  one  of  the  illuminating  constit- 
uents of  coal  gas.  When  ethylene  burns,  the  complete 
combustion  is  represented  thus  — 

C2H4     +     302     =         2C02         +     2H20 
Ethylene  Carbon  Dioxide  Water 

If  mixed  with  oxygen  in  this  proportion  and  ignited,  the 
mixture  explodes. 

Other  numbers  of  this  series  are  propylene  (C3H6)  and  butylene 
(C4H8).  These  are  unsaturated  hydrocarbons.  Unlike  the  paraffins, 
they  form  addition  products  by  uniting  directly  with  other  substances, 
especially  chlorine,  thus  — 

C2H4     +     C12     =  C2H4C12 

Ethylene  Ethylene  Chloride 

Ethylene  chloride  is  one  of  the  two  dichlorethanes ;  they  have  the 
same  percentage  composition,  molecular  weight,  and  formula  (C2H4C12), 
but  are  very  different  compounds.  They  illustrate  isomerism  and  are 
called  isomers.  This  kind  of  isomerism  is  called  metamerism.  The 
difference  in  properties  is  believed  to  be  due  to  a  different  arrangement 
of  the  atoms  in  the  molecules.  Isomerism  occurs  frequently  among 
carbon  compounds. 


Acetylene.  205 

Acetylene  is  formed  by  the  direct  union  of  hydrogen 
and  carbon  when  an  electric  arc  is  produced  between  two 
carbon  rods  in  hydrogen  gas.  This  method  of  formation, 
though  not  convenient,  is  interesting,  because  no  other  hy- 
drocarbon has  as  yet  been  directly  built  up  from  its  elements. 
A  small  quantity  is  present  in  coal  gas.  It  is  also  formed 
by  the  incomplete  combustioft  of  coal  gas,  e.g.  when  the 
flame  of  a  Bunsen  burner  strikes  back  and  burns  at  the 
base  (see  Bunsen  Burner).  Acetylene  is  now  prepared 
cheaply  on  a  large  scale  by  treating  calcium  carbide  with 
water,  thus  — 

CaC2       +     2H2O     =     C2H2     +     Ca(OH)2 

Calcium  Carbide  Acetylene 

Acetylene  is  a  colorless  gas,  and,  if  impure,  has  an  offen- 
sive odor.  It  is  poisonous  if  breathed  in  large  quantities, 
but  much  less  dangerous  than  gases  containing  carbon 
monoxide.  It  is  lighter  than  air,  its  density  being  about 
0.92.  Water  at  the  ordinary  temperature  dissolves  its  own 
volume  of  the  gas.  Reliable  tests  show  that  acetylene 
does  not  act  upon  any  common  metal  or  alloy,  though  it 
forms  explosive  compounds  with  salts  of  metals,  especially 
copper.  As  a  precaution,  copper  and  brass  are  seldom 
used  in  large  vessels  containing  or  generating  acetylene, 
though  they  might  be  safely  used  on  small  vessels  like 
bicycle  lamps. 

Under  a  pressure  of  40  atmospheres  and  a  temperature  of  20°  C.  it 
liquefies.  Cylinders  of  liquid  acetylene  have  exploded,  causing  loss  of 
life  and  destruction  of  property,  and  its  use  in  this  form  has  been  pro- 
hibited in  some  localities.  Under  ordinary  atmospheric  conditions  acety- 
lene will  not  explode.  If  compressed,  it  will  explode  when  a  spark  or 
flame  is  brought  near  it.  A  mixture  of  acetylene  and  air,  if  ignited, 
explodes.  The  mixture  to  be  explosive,  however,  must  contain  from 
about  3  to  65  per  cent  of  acetylene  (a  condition  hardly  possible 


206 


Descriptive  Chemistry. 


except  from  sheer  carelessness),  because  the  disagreeable  odor  reveals 
the  presence  of  the  gas.  Acetylene  must  be  used  with  the  same  precau- 
tion as  any  other  illuminating  gas. 

Acetylene  is  found  by  analysis  to  contain  only  carbon  and  hydrogen 
combined  in  the  ratio  of  12  to  I  by  weight.  Its  vapor  density  is  13. 
Therefore  its  molecular  weight  must  be  26  and  its  formula  C2H.,. 

Acetylene  is  an  unsaturated  hydrocarbon,  and  like  ethylene  combines 
directly  with  bromine,  hydrogen,  and  other  elements.  When  passed 
into  silver  or  copper  solutions,  it  forms  explosive  compounds  called 
acetylides  (e.g.  Ag2C2  and  Cu2C2).  Heated  to  a  high  temperature,  it 
changes  into  other  hydrocarbons,  one  being  benzene,  thus  — 

3C2H2       =       C6H6 
Acetylene  Benzene 

At  a  very  high  temperature  (about  800°  C.)  it  decomposes  into  carbon 
and  hydrogen.  The  change  of  acetylene  into  benzene  illustrates  po- 
lymerism.  Polymers  have  the  same  percentage  composition,  but 
different  molecular  weights  (see  Isomerism). 

Acetylene  as  an  Illuminant.  —  Acetylene  burns  in  the 
air  with  a  luminous,  smoky  flame.  But  when  air  is  mixed 

with  the  gas  as  the  latter 
issues  from  a  small  opening, 
the  mixture  burns  with  a 
brilliant,  white  flame,  which 
does  not  smoke.  It  is  grad- 
ually coming  into  use  as  an 
illuminant.  The  flame  is 
almost  like  sunlight,  hence 
by  the  acetylene  flame  most 
colors  appear  the  same  as  in 
daylight.  It  is  also  adapted 
for  taking  photographs,  since 
its  action  closely  resembles 
that  of  the  sun.  It  is  a  diffusive  light,  and  the  flame  is 
much  smaller  than  an  ordinary  gas  flame  of  the  same 
lighting  power  (Fig.  32). 


FlG.  32.  —  Relative  size  of  acetylene 
and  illuminating  gas  flames  giving  the 
same  amount  of  light.  The  acetylene 
(smaller)  flame  consumes  only  one 
tenth  as  much  gas  an  hour  as  the  illu- 
minating gas  flame.  (One  half  actual 
size.) 


Petroleum.  207 

With  a  proper  burner  the  combustion  of  acetylene  is 
complete,  and  may  be  represented  thus  — 

2C2H2     +     5O2     =         4CO2         +     2H2O 
Acetylene  Oxygen         Carbon  Dioxide  Water 

In  most  acetylene  burners  the  gas  issues  from  two 
small  holes  drilled  at  an  angle,  so  that  the  jets  strike 
each  other  and  produce  a  flat  flame 
(Fig.    33).       Other     holes,    properly 
located,   permit   air   to    be   drawn   in 
mechanically  by  the   acetylene   as   it 
rushes  through  the  burner.     The  open- 
ings for  the  mixture  are  so  fine  that 
FIG.  33.  — Acety-    the  flame  cannot  strike  back  and  cause    FlG  34._  Acety- 
lene flame.         an  explosion  (Fig.  34).  lene  burner. 

Generation  of  Acetylene.  —  The  ease  with  which  acetylene  is  gener- 
ated can  be  shown  by  putting  a  little  water  in  a  test  tube  and  then  drop- 
ping in  small  lumps  of  calcium  carbide.  The  gas  bubbles  through  the 
liquid ;  after  the  action  has  proceeded  long  enough  to  expel  the  air, 
the  acetylene  may  be  lighted  by  holding  a  burning  match  at  the  mouth 
of  the  tube.  On  a  larger  scale,  the  gas  can  be  generated  by  putting  the 
calcium  carbide  into  a  flask  provided  with  a  dropping  funnel  and  de- 
livery tube,  and  allowing  water  to  drop  slowly  upon  the  carbide ;  the 
gas  thus  generated  can  be  collected  in  bottles  over  water.  There  are 
two  classes  of  commercial  generators.  In  one,  water  is  added  to  the 
calcium  carbide,  but  in  the  other  the  carbide  drops  into  the  water.  The 
intense  heat  liberated  when  calcium  carbide  interacts  with  water  de- 
composes acetylene ;  hence,  a  generator  to  be  effective  and  safe  should 
be  constructed  so  that  this  heat  will  be  absorbed.  The  first  class  of 
generators  is  dangerous,  except  when  a  small  quantity  of  gas  is  desired, 
as  on  the  lecture  table  or  in  a  bicycle  lantern.  -In  the  second  class,  a 
small  amount  of  calcium  carbide  drops  automatically  into  a  large  vol- 
ume of  water  as  fast  as  the  gas  is  needed,  thus  insuring  a  pure,  cool 
gas,  and  eliminating  the  danger  of  an  explosion.  A  pound  of  calcium 
carbide  yields  about  five  cubic  feet  of  acetylene  gas. 

Petroleum  is  the  source  of  many  useful  hydrocarbons. 
It  is  an  oily  liquid  obtained  from  the  earth  in  many  parts 
of  the  world.  In  the  United  States  the  chief  localities  are 


208  Descriptive  Chemistry. 

Ohio,  New  York,  Pennsylvania,  West  Virginia,  Kentucky, 
Indiana,  Colorado,  Texas,  and  California.  The  immense 
deposits  in  Russia  are  in  the  Baku  district  on  the  Caspian 
Sea.  Some  is  also  found  in  Canada,  India,  Japan,  and 
Austria. 

Crude  petroleum  is  a  thick  liquid,  with  an  unpleasant 
odor.  Its  color  varies  from  straw  to  greenish  black,  and 
most  kinds  are  greenish  in  reflected  light.  It  usually  floats 
upon  water.  Its  composition  is  complex,  but  all  varieties 
are  essentially  mixtures  of  many  hydrocarbons.  Ameri- 
can oils  contain  chiefly  members  of  the  paraffin  series. 
Some  varieties  contain  compounds  of  nitrogen  and  of 
sulphur. 

In  some  localities  the  oil  issues  from  the  earth,  but  it  is  usually  neces- 
sary to  drill  through  rocks  and  insert  a  pipe  into  the  porous  rock 
containing  oil.  At  first  the  oil  often  "shoots'1  out  of  the  well  in 
tremendous  volumes,  owing  to  the  pressure  of  the  confined  gas,  but 
after  a  time  a  pump  is  needed  to  draw  it  to  the  surface.  The  oil  is  then 
forced  by  powerful  pumps  through  large  pipes  to  central  points  for 
storage  or  for  delivery  to  refineries,  which  are  often  many  miles  from 
the  oil  well.  This  network  of  pipes  in  the  eastern  United  States  is  over 
25,000  miles  long. 

Some  crude  petroleum  is  used  in  making  water  gas  (see 
below),  and  as  fuel  on  locomotives  and  steamships,  but 
most  of  it  is  separated  into  various  commercial  products. 
This  process,  which  also  involves  purification,  is  called  re- 
fining. The  petroleum  is  distilled  in  huge  iron  vessels, 
and  the  vapors  are  condensed  as  they  pass  through  coiled 
pipes  immersed  in  cold  water.  Certain  products  are  ob- 
tained from  the  residue  left  in  the  still. 

The  different  distillates,  which  are  collected  in  separate  tanks,  are 
further  separated  and  purified  by  redistillation.  The  commercial 
products  obtained  from  the  first  distillation  are  cymogene,  rhigolene, 
gasolene,  naphtha,  benzine,  and  kerosene.  These  liquids  are  mixtures 


Natural  Gas.  209 

of  several  different  hydrocarbons.  They  are  widely  used  as  solvents, 
fuels,  and  in  making  gas. 

Kerosene  is  the  well-known  illuminating  oil.  Being  the  most  valu- 
able product  from  petroleum,  it  is  very  carefully  freed  from  inflammable 
liquids  and  gases,  which  might  cause  an  explosion,  and  from  tarry 
matter  and  semi-solid  hydrocarbons,  which  would  clog  the  wicks  of 
lamps.  This  is  done  by  agitating  it  successively  with  sulphuric  acid, 
sodium  hydroxide,  and  water.  Commercial  kerosene  must  have  a  legal 
flashing  point.  This  is  "the  temperature  at  which  the  oil  gives  off 
sufficient  vapor  to  form  a  momentary  flash  when  a  small  flame  is 
brought  near  its  surface."  In  most  states  the  flashing  point  is  44°  C. 
(or  iii°F.). 

From  the  residuum  left  in  the  still  after  the  first  distillation  many 
grades  of  lubricating  oil,  vaseline,  and  paraffin  wax  are  obtained 
by  further  treatment.  Mineral  lubricating  oils  have  largely  replaced 
animal  and  vegetable  oils.  Vaseline  finds  extensive  use  as  an  ointment. 
Paraffin  wax  is  used  to  make  candles,  to  water-proof  paper,  to  extract 
oils  from  plants  and  flowers,  and  as  a  coating  for  many  substances, 
thereby  producing  a  smooth  surface  or  facilitating  slow  combustion  (as 
in  parlor  matches).  The  final  residue  is  coke.  Hydrocarbons  are 
often  extracted  from  it,  some  is  made  into  electric  light  carbons,  and 
some  is  used  as  a  fuel. 

This  vast  industry  yields  over  two  hundred  different  commercial 
products,  many  of  them  being  indispensable  to  the  comfort  and  conven- 
ience of  mankind.  In  1901  the  United  States  produced  over  69,000,000 
barrels  of  crude  petroleum. 

The  Origin  of  Petroleum  is  doubtful.  Some  think  it  was  produced 
by  the  decomposition  or  slow  distillation  of  plants  and  animals. 
Recently  it  has  been  suggested  that  it  resulted  from  the  interaction  of 
water  and  metallic  carbides,  especially  iron  carbide,  at  great  depths. 

Natural  Gas  is  a  combustible  gas,  which  issues  from  the 
earth  in  many  places.  Methane  is  the  principal  constituent 
of  the  mixture.  It  is  used  as  a  fuel  for  heating  houses, 
generating  steam,  and  manufacturing  iron,  steel,  glass, 
brick,  and  pottery. 

In  Ohio,  Indiana,  and  other  gas-producing  regions  of  the  United 
States,  wells,  like  petroleum  wells,  are  drilled  for  the  escape  of  natural 


2io  Descriptive  Chemistry. 

gas,  which  is  distributed  to  consumers  through  pipes  similar  to  those 
used  for  illuminating  gas.  Enormous  quantities  are  consumed  in  the 
United  States,  the  annual  product  being  valued  at  over  $20,000,000. 

Illuminating  Gas.  —  Besides  acetylene  there  are  other 
kinds  of  illuminating  gas.  Coal  gas  and  water  gas  are  the 
most  common. 

Coal  Gas  is  made  by  distilling  bituminous  coal  and  puri- 
fying the  volatile  product.  The  hydrogen  in  the  coal 
passes  off  partly  as  free  hydrogen,  and  partly  in  combina- 
tion with  carbon  as  hydrocarbons,  and  with  nitrogen  as 
ammonia.  The  ammonia,  carbon  dioxide,  and  sulphur 
compounds  are  regarded  as  impurities,  and  are  removed 
before  the  gas  is  sent  to  the  consumer.  The  essential 
parts  of  a  coal-gas  plant  are  shown  in  Figure  35. 

The  coal  is  distilled  in  a -shaped  retorts,  made  of  fire  clay  and 
about  eight  feet  long.  Six  or  more  retorts  are  arranged  in  tiers  form- 
ing a  group  or  bench,  so  that  all  the  retorts  of  a  bench  can  be  heated 
by  a  single  fire  —  usually  of  coke.  Several  benches  placed  end  to  end 
constitute  a  stack.  The  retorts  are  heated  red  hot,  and  about  two  hun- 
dred pounds  of  coal  are  evenly  distributed  on  the  bottom  of  each  retort 
with  a  long  iron  scoop,  and  the  mouth  is  quickly  and  tightly  closed  by 
an  iron  lid.  The  distillation  continues  from  four  to  six  hours,  during 
which  the  temperature  often  reaches  1200°  C.  The  lid  is  then  removed, 
the  red-hot  coke  is  pushed  or  raked  out,  and  another  charge  of  coal  is 
quickly  introduced.  The  coke  is  quenched  with  water  to  prevent  fur- 
ther combustion.  Some  of  it  is  used  for  heating  the  retorts,  but  a  part 
is  sold. 

The  volatile  products  pass  from  each  retort  up  through  a  standpipe, 
down  the  dip  pipe,  and  bubble  through  water  into  the  hydraulic  main. 
This  is  a  horizontal,  half-round  pipe  extending  the  whole  length  of  the 
stack.  Here  some  of  the  tar  is  deposited  and  ammonium  compounds 
are  dissolved  by  the  water  which  flows  constantly  through  the  main. 
This  water  is  kept  at  the  same  level  and  acts  as  a  "  seal "  to  prevent  the 
gas  from  passing  back  into  the  retorts.  The  ammoniacal  liquor  and 
tar  flow  into  a  tar  well. 

From  the  hydraulic  main  the  gas  which  is  hot  and  impure  passes 


Illuminating  Gas. 


211 


SJ.UOJ.3U 


212  Descriptive  Chemistry. 

into  the  condenser.  This  is  a  series  of  vertical  iron  pipes,  several 
hundred  feet  long.  They  are  connected  at  the  top,  but  they  open  at 
the  bottom  into  a  series  of  boxes  so  constructed  that  the  gas  must  pass 
through  the  entire  length  of  the  pipes,  while  the  tar  and  ammoniacal 
liquor  flow  into  'the  tar  well.  The  main  object  of  the  condenser  is  to 
cool  the  gas  slowly  and  condense  and  remove  the  tar. 

An  exhauster,  in  most  plants,  draws  or  forces  the  gas  from  the 
hydraulic  main  through  the  condenser  into  the  scrubber  and  onward 
through  the  purifiers  into  the  gas  holder.  The  exhauster  also  reduces 
the  pressure  in  the  retorts  and  regulates  the  pressure  in  the  holder  (see 
below) . 

The  scrubber  is  a  washing  machine.  Its  purpose  is  to  remove  the 
remaining  ammonia,  part  of  the  carbon  dioxide,  and  hydrogen  sulphide 
gas,  and  the  last  traces  of  tar.  Scrubbers  vary  in  construction.  One 
form  is  a  double  tower  filled  with  wooden  slats  or  with  trays  covered 
with  coke  or  pebbles  over  which  ammoniacal  liquor  slowly  trickles  in 
the  first  part  and  pure  water  in  the  second.  The  gas  enters  at  the 
bottom,  meets  the  descending  liquid,  and  is  thoroughly  washed. 
Another  form  widely  used  consists  of  a  cylindrical  vessel  in  which 
numerous  wooden  slats  revolve  in  compartments  and  dip  into  am- 
moniacal liquor  or  water  at  the  bottom.  The  liquid  forms  a  film  on 
the  slats  and  absorbs  the  ammonia  and  other  gases,  while  the  resulting 
solution  mixes  with  liquor  at  the  bottom  and  flows  into  the  proper  well. 
Sometimes  a  separate  tar  extractor  is  connected  with  the  scrubber. 
This  is  a  tower  filled  with  perforated  plates,  which  catch  and  remove 

the   tar   mechanically  as    the   gas 
passes  through  into  the  scrubber. 

From    the     scrubber    the    gas 
passes   into  the  purifiers.     Their 
FIG.  36. -Slat  frame  (or  grid)  used  in    chief    purpose    is    to    remove    the 
the  lime  purifier.  remaining  carbon  dioxide  and  sul- 

phur compounds.  They  are  shal- 
low, rectangular  iron  boxes  provided  with  slat  frames  loosely  covered 
with  lime  (Fig.  36).  In  some  plants  iron  oxide  is  used  as  the  purifying 
material. 

The  purified  gas  next  passes  through  a  large  meter,  which  records 
its  volume,  into  a  gas  holder.  The  holder  is  an  enormous,  cylindrical, 
iron  tank  in  which  the  gas  is  stored.  It  floats  in  a  cistern  of  water,  and 
rises  or  falls  as  the  gas  enters  or  leaves.  Weights  and  the  pressure 


Water  Gas.  213 

from  the  exhauster  so  balance  it  that  it  exerts  just  enough  pressure  to 
force  the  gas  through  the  pipes  to  the  consumer. 

A  ton  of  good  coal  yields  about  10,000  cubic  feet  of  gas,  1400  pounds 
of  coke,  120  pounds  of  tar,  20  gallons  of  ammoniacal  liquor,  and  a  vary- 
ing amount  of  gas  carbon.  The  coke  is  a  valuable  fuel  and  finds  a 
ready  sale.  The  tar,  or  coal  tar  as  it  is  often  called,  collected  from  the 
hydraulic  main  and  condenser,  is  a  thick,  black,  foul-smelling  liquid. 
It  was  formerly  thrown  away.  Some  is  used  for  preserving  timber, 
making  tarred  paper  and  concrete,  and  as  a  protective  paint.  Most  of 
it  is  now  separated  by  distillation  into  its  more  important  constituents, 
especially  benzene  (C6HC) .  These  carbon  compounds  and  their  numer- 
ous derivatives  appear  in  commerce  as  oils,  medicines,  dyestufFs,  flavors, 
perfumes,  and  other  useful  products.  The  ammoniacal  liquor  from  the 
hydraulic  main,  condenser,  and  scrubber  is  the  source  of  ammonia  and 
its  compounds.  Gas  carbon  is  the  hard  deposit  which  collects  on  the 
inside  of  the  retort,  and  is  used  in  the  electrical  industries  (see  Gas 
Carbon).  The  sale  of  these  by-products  reduces  the  cost  of  making 
the  coal  gas. 

Water  Gas  is  made  by  forcing  steam  through  a  mass  of 
red-hot  coal  and  mixing  the  gaseous  product  with  hot  gases 
obtained  from  oil.  The  essential  parts  of  the  apparatus 
are  shown  in  Figure  37. 

Air  is  forced  through  the  coal  fire  in  the  generator,  and  the  hot 
gases  which  are  produced  pass  down  the  carburetor,  up  into  the  super- 
heater, and  escape  through  its  top  into  the  open  air.  This  operation 
lasts  about  four  minutes,  and  is  called  the  "  blow."  It  heats  the  fire 
brick  inside  the  carburetor  and  superheater  intensely  hot,  air  often  being 
forced  in  to  raise  the  temperature.  The  air  valves  and  the  top  of  the 
superheater  are  now  closed,  and  the  "  run "  begins,  which  lasts  about 
six  minutes.  Steam  is  forced  into  the  generator  at  the  bottom.  In 
passing  through  the  mass  of  incandescent  carbon  the  steam  and  carbon 
interact  thus  — 

C          +        H2O  CO  +  H2 

Carbon  Steam       Carbon  Monoxide  Hydrogen 

This  mixture  of  hydrogen  and  carbon  monoxide  burns  with  a  feeble 
flame,  and  before  it  can  be  used  as  an  illuminating  gas  it  must  be 


214 


Descriptive  Chemistry. 


Characteristics  of  Illuminating  Gases.        215 

enriched  with  gases  which  are  illuminants.  Therefore,  the  mixed  gases 
pass  to  the  top  of  the  carburetor,  where  they  meet  a  spray  of  oil.  And 
as  the  gaseous  mixture  passes  down  the  carburetor  and  up  the  super- 
heater, the  hydrocarbons  of  the  oil  are  transformed  by  the  intense  heat 
into  hydrocarbons  that  do  not  liquefy  when  the  gas  is  cooled.  The  ad- 
dition of  hydrocarbons  is  called  carbureting.  From  the  superheater 
the  water  gas  passes  through  the  purifying  apparatus  into  a  holder. 

Water  gas  is  seldom  burned  alone,  but  is  usually  mixed 
with  60  or  70  per  cent  of  coal  gas.  This  mixture  is  popu- 
larly called  "  illuminating  gas."  Owing  to  the  high  percen- 
tage of  carbon  monoxide,  water  gas  and  gases  containing 
it  are  poisonous. 

Characteristics  of  Illuminating  Gases.  —  Both  coal  gas 
and  water  gas  have  a  disagreeable  odor.  They  are  mix- 
tures having  a  composition  which  varies  with  the  coal 
used,  the  temperature  reached,  and  the  degree  of  purifica- 
tion attained.  The  following  table  shows  the  average  — 


COMPOSITION  OF  ILLUMINATING  GASES. 


CONSTITUENTS. 

COAL  GAS. 

WATER  GAS. 

Marsh  2fas     .     .          

•JA    C 

19  8 

Ethylene  (and  other  illuminants) 
Hydrogen                

J^fO 
5.0 
AQ  O 

16.6 

•52     I 

Carbon  monoxide            ... 

7  2 

J^.l 

26  I 

Carbon  dioxide 

I    i 

30 

•5.2 

•M 

2.4. 

Both  kinds  of  illuminating  gas  may  contain  a  little  oxygen,  and 
traces  of  ammonia  and  hydrogen  sulphide  gases.  Nitrogen  and  the 
last  portions  of  carbon  dioxide  are  impurities  not  easily  removed. 
Marsh  gas,  hydrogen,  and  carbon  monoxide  burn  with  a  feeble  (non- 
yellow)  flame,  and  are  often  called  diluents ;  they  furnish  heat,  but  no 
light. 


216  Descriptive  Chemistry. 

The  luminosity  of  illuminating  gas  depends  mainly 
upon  the  presence  of  hydrocarbons  containing  a  relatively 
large  proportion  of  carbon.  Acetylene  gas,  which  gives 
such  a  brilliant  light,  consists  almost  wholly  of  this  hydro- 
carbon containing  90  per  cent  of  carbon.  The  most  im- 
portant illuminants  in  coal  gas  and  water  gas  are  ethylene 
and  similar  hydrocarbons,  acetylene,  and  benzene  (C6H6). 

The  commercial  value  of  an  illuminating  gas  depends  upon  its  illu- 
minating power.  This  property  is  measured  by  a  photometer  and  is 
expressed  in  •*  candles."  The  determination  is  made  by  comparing  the 
light  produced  by  burning  the  gas  in  a  standard  burner  at  the  rate  of 
five  cubic  feet  an  hour  with  the  light  produced  by  a  standard  wax  candle 
burning  at  the  rate  of  120  grains  (7.77  gm.)  an  hour.  If  the  gas  flame 
is  20  times  brighter  than  the  candle  flame,  then  the  candle  power  of  the 
gas  is  20.  The  candle  power  of  ordinary  coal  gas  is  about  17,  and 
that  of  water  gas  is  about  25.  Ordinary  illuminating  gas  has  a  candle 
power  of  about  20,  since  it  is  usually  a  mixture  of  coal  gas  and  water 
gas. 

Flame.  — A  flame  is  a  mass  of  burning  gas.  Ordinarily 
it  is  gas  combining  chemically  with  the  oxygen  of  the  air. 
In  the  illuminating  gas  flame  the  gas  itself  is  burning  in 
the  air.  In  a  lamp  flame  the  gas  which  burns  comes  from 
the  oil  which  is  drawn  up  the  wick  by  capillary  attraction, 
and  then  volatilized  by  the  heat.  Similarly,  in  a  candle 
flame  the  burning  gas  comes  from  the  melted  wax.  The 
flame  produced  by  most  burning  hydrocarbons  is  yellowish 
white. 

The  hydrocarbon  flame  has  several  distinct  parts,  though 
the  structure  of  the  flame  is  essentially  the  same,  whether 
produced  by  burning  illuminating  gas,  kerosene  oil,  or  can- 
dle wax.  The  candle  flame  may  be  taken  as  the  type.  An 
examination  of  the  enlarged  vertical  section  shown  in  Fig- 
ure 38  reveals  four  somewhat  conical  portions,  (i)  Around 
the  wick  there  is  a  black  cone  (A),  filled  with  combustible 


Flame. 


217 


FIG.  38.  —  Candle 
flame. 


gases  formed  from  the  melted  wax.  They  do  not  burn  be- 
cause no  oxygen  is  present.  With  a  glass  tube  of  fine 
bore  it  is  possible  to  draw  off  these  gases 
from  a  large  flame  and  light  them  at  the 
upper  end  of  the  tube.  (2)  Around  the 
lower  part  of  the  dark  cone  is  a  faint  bluish 
cup-shaped  part  (£>,  B).  It  is  the  lower  por- 
tion of  the  exterior  cone  where  complete 
combustion  of  the  gases  occurs,  since  plenty 
of  oxygen  from  the  air  reaches  this  portion. 
(3)  Above  the  dark  cone  is  the  luminous 
portion  (C).  It  is  the  largest  and  most 
important  part  of  the  flame.  It  is  popu- 
larly spoken  of  as  "  the  flame."  Combus- 
tion is  incomplete  here,  because  little  or  no 
oxygen  can  pass  through  the  exterior  cone.  The  tempera- 
ture is  high,  however,  and  the  hydrocarbons  undergo 
complex  changes.  Acetylene  is  probably  formed.  The 
most  characteristic  change  is  the  liberation  of  small  par- 
ticles of  carbon.  This  liberated  carbon  heated  to  incan- 
descence by  the  burning  gases  makes  the  flame  luminous. 
The  carbon  glows  but  does  not  burn  up,  because  little  or 
no  oxygen  is  present.  A  crayon  or  glass  rod  held  in  this 
part  of  the  flame  is  at  once  coated  with  soot,  which  consists 
of  fine  particles  of  carbon.  The  exterior  cone  (D,  D)  is 
almost  invisible.  Here  combustion  is  complete,  because 
the  oxygen  of  the  air  changes  all  the  carbon  into  carbon 
dioxide.  That  this  is  the  hottest  region 
may  be  easily  shown  by  pressing  a  piece  of 
stiff  white  paper  for  an  instant  down  upon 
the  flame  almost  to  the  wick.  The  paper 
FIG.  39.- Paper  wi^  frQ  charred  by  the  outer  part  of  the 

charred  by  a  can- 
dle flame.  flame,  as  shown  in  Figure  39. 


2i 8  Descriptive  Chemistry. 

These  four  portions  may  be  found  in  all  luminous  hydrocarbon 
flames,  whatever  the  shape.  An  ordinary  gas  flame  is  flattened  by  forc- 
ing the  gas  flame  through  a  narrow  slit  in  the  burner,  so  that  the  flame 
will  give  more  light.  The  blue  part  is  easily  seen,  however,  when  the 
gas  flame  is  turned  low  or  looked  at  through  a  small  opening ;  the  dark 
and  yellow  parts  are  always  visible  —  the  latter  being  intentionally  en- 
larged. The  flat  or  circular  flame  of  an  oil  lamp  likewise  presents  the 
same  characteristics. 

The  gaseous  products  of  the  combustion  of  hydrocarbons 
are  water  vapor  and  carbon  dioxide.  A  bottle  in  which  a 
candle  is  burning  has,  at  first,  a  deposit  of  moisture  on  the 
inside ;  and  if  the  candle  is  removed  and  limewater  added, 
the  presence  of  carbon  dioxide  is  shown  by  the  milkiness  of 
the  limewater.  The  oxygen  needed  by  the  burning  hydro- 
carbons is  obtained  from  the  air.  If  not  enough  oxygen  is 
present,  the  flame  smokes,  i.e.  the  carbon  is  thrown  off  into 
the  air  before  the  particles  are  heated  hot  enough  to  glow. 
All  oil  lamps  are  so  constructed  that  air  enters  the  burner 
below  the  flame.  Large  oil  lamps  have  a  central  opening 
through  which  a  large  volume  of  air  passes  up  inside  the 
circular  flame.  Otherwise  the  lamp  would  burn  with  a 
very  smoky  flame. 

The  luminosity  of  hydrocarbon  flames  is  affected  by  other  things 
besides  the  presence  of  glowing  carbon.  One  of  these  is  temperature. 
Gases  cooled  before  being  burned  give  poor  light.  A  candle  flame  may 
be  cooled  enough  to  extinguish  it.  Thus,  if  a  coil  of  copper  wire  is 
lowered  upon  a  candle  flame,  the  flame  smokes,  loses  its  yellow  color, 
and  finally  goes  out ;  but  if  a  coil  of  hot  wire  is  used,  the  flame  burns 
unchanged.  Gases,  as  well  as  solids  and  liquids,  have  a  kindling  tem- 
perature, i.e.  a  temperature  to  which  they  must  be  heated  before  they 
"  catch  fire."  This  temperature  differs  with  different  substances.  As 
we  lower  the  temperature  of  gases  burning  with  a  luminous  flame,  their 
luminosity  decreases,  and  below  their  kindling  point  they  will  not  burn. 
The  density  of  the  gases  in  the  flame  and  of  the  atmosphere  itself  like- 
wise modifies  luminosity.  A  candle  flame  was  found  by  experiment  to 
be  smaller  on  the  top  of  Mont  Blanc  than  at  the  base. 


The  Bunsen  Burner  and  its  Flame.         219 


O 


Not  all  flames  are  luminous.  The  hydrogen  flame  is  almost  invisible, 
and  the  flames  of  carbon  monoxide  and  methane  are  a  faint  blue.  These 
flames  yield  no  solid  particles  of  carbon,  but  only  gaseous  products.  The 
most  common  non-luminous  flame  is  the  Bunsen  flame. 

The  Bunsen  Burner  and  its  Flame.  —  When  illuminat- 
ing gas  is  mixed  with  air  before  burning,  and  the  mixture 
burned  in  a  suitable  burner,  a  flame  is  produced  which  is 
non-luminous  -and  very  hot.  The 
temperature  of  the  hottest  part  is 
about  1 500°  C.  This  flame  deposits 
no  carbon,  since  its  products  are 
entirely  gaseous.  Such  a  flame  is 
called  the  Bunsen  flame,  because 
it  is  produced  in  a  burner  devised 
by  the  German  chemist  Bunsen. 
This  burner  is  constantly  used  in 
chemical  laboratories  as  a  source 
of  heat,  and  modified  forms  have 
numerous  uses.  One  form,  for 
example,  furnishes  the  heat  in  the 
gas  range  used  for  cooking.  The 
parts  of  an  ordinary  Bunsen  burner 
are  shown  in  Figure  40.  The  gas 
enters  the  base  and  escapes  through 
a  very  small  opening  into  the  long 
tube,  which  screws  down  upon  this 
opening.  At  the  lower  end  of  the 
long  tube  there  are  two  holes, 
through  which  air  is  drawn  by  the  gas  as  it  rushes  out  of 
the  small  opening.  The  gas  and  air  mix  as  they  rise  in  the 
tube,  and  this  mixture  of  air  and  gas  burns  at  the  top  of 
the  long  tube.  The  size  of  the  air  holes  at  the  bottom  of 
the  long  tube  may  be  changed  by  a  movable  ring,  thus 


FIG.  40.  —  Parts  of   a  Bunsen 
burner. 


220  Descriptive  Chemistry. 

varying  the  volume  of  the  entering  air.  When  the  holes 
are  open,  the  typical  colorless,  hot  Bunsen  flame  is  formed. 
The  combustion  of  the  hydrocarbons  is  practically  com- 
plete. They  burn  up  before  particles  of  carbon  are 
liberated,  thus  making  the  flame  non-luminous  and  free 
from  soot.  Apparatus  heated  by  this  flame  is  not  black- 
ened. The  Bunsen  flame  may  be  made  momentarily 
luminous  by  shaking  or  blowing  fine  particles  into  the 
flame,  —  such  as  powdered  charcoal  dust,  finely  divided 
metals,  and  sodium  compounds. 

It  was  formerly  believed  that  the  non-luminous  character  of  the 
Bunsen  flame  is  solely  due  to  the  complete  combustion  of  the  carbon  by 
the  oxygen  of  the  entering  air.  Recent  experiments  have  shown,  how- 
ever, that  the  result  is  partly  due  to  the  diluting  action  of  the  nitrogen*— 
The  gas  burns  at  top  of  the  tube  and  not  inside,  because  the  proper 
mixture  of  gas  and  air  flows  out  more  quickly  than  the  flame  can  travel 
back.  If  the  gas  supply  is  slowly  decreased,  the  flame  becomes  smaller 
and  finally  disappears  with  a  slight  explosion.  This  change  is  called 
"striking  back."  It  is  due  to  the  fact  that  the  tube  contains  an  explo- 
sive mixture  of  air  and  illuminating  gas,  through  which  the  flame  travels 
faster  than  the  mixture  escapes  from  the  tube.  This  explosion  illus- 
trates in  a  small  way  what  often  happens  when  a  mixture  of  air  and 
illuminating  gas  is  ignited.  Sometimes  the  flame  is  not  extinguished, 
but  burns  within  (and  sometimes  without)  the  tube.  This  flame  has  a 
pale  color,  a  disagreeable  odor,  and  deposits  soot. 

The  Bunsen  flame  has  many  characteristic  properties. 
Its  color  is  bluish,  and  the  different  corres  have  different 
colors.  There  are  really  three  cones:  (i)  the  blue  or 
greenish  inner  one  of  unburned  gases ;  (2)  the  very  faint 
blue  middle  one  ;  (3)  and  the  outer  one,  which  is  pale  blue, 
and  represents  the  blue  cone  in  the  candle  flame.  The 
middle  and  outer  cones  are  not  always  easily  distinguished ; 
and  for  all  practical  purposes  it  is  convenient  to  divide  the 
flame  into  two  parts,  —  an  inner  cone  of  unburned  gases 


Oxidizing  and  Reducing  Flames.  221 


FIG.  41.  —  The  effects  of  wire  gauze  on  a 
Bunsen  flame. 


and  an  outer  cone  in  which  all  the  carbon  is  consumed. 
Combustible  gases  may  be  drawn  off  by  a  tube  from  the 
inner  cone  and  ignited.  A  match  laid  for  an  instant  across 
the  top  of  the  tube  is 
charred  only  at  the  two 
points  where  it  touches 
the  outer  cone ;  and  a 
sulphur  match'-  suspended 
by  a  pin  across  the  top 
of  an  unlighted  burner 
is  not  kindled  when  the 
gas  is  first  lighted.  A 
piece  of  wire  gauze  pressed  down  upon  the  flame  shows 
a  dark  central  portion  surrounded  by  a  luminous  ring. 
The  flame  is  beneath  the  gauze,  although  the  gas  passes 
freely  through  it  and  escapes.  If  the  gas  is  extinguished 
and  then  relighted  above  the  gauze,  it  will  burn  above 
but  not  beneath  (Fig.  41).  The  gauze  cools  the  gas  below 
its  kindling  temperature. 

The  minerfs  safety  lamp  invented  by  Davy  depends 
upon  this  last  principle.  It  is  an  oil  lamp  surrounded 
by  a  cylinder  of  fine  wire  gauze  (Fig.  42).  When 
taken  into  a  mine  where  there  are  explosive  gases  (fire 
damp),  the  flame  continues  to  burn  inside,  though  its 
size  and  color  change.  The  gas  often  enters  the  lamp 
and  burns  inside,  but  the  flame  within  does  not  ignite 
the  gases  without  because  the  wire  gauze  keeps  them 
cooled  below  their  kindling  temperature.  Hence  an 
explosion  is  often  prevented.  When  miners  notice 
changes  in  the  lamp  flame,  they  usually  seek  a  safe 
FIG.  42.  —  One  place, 
form  of  Davy's  * 

safety  lamp.  Oxidizing  and  Reducing  Flames.  — The 

outer  portiqn  of  the  Bunsen  flame  is  called  the  oxidizing 
flame,  because  here  the  oxygen  is  freely  given  to  sub- 


222 


Descriptive  Chemistry. 


-~Y—- A 


stances.  The  inner  portion  is  called  the  reducing  flame, 
because  here  the  hydrocarbons  withdraw 
oxygen.  A  sketch  of  the  general  relation 
of  these  flames  is  shown  in  Figure  43.  A 
is  the  most  effective  part  of  the  oxidizing 
flame,  and  B  of  the  reducing  flame.  At 
A  metals  are  oxidized,  and  at  B  oxygen 
compounds  are  reduced. 

Sometimes  a  long  tube  with  a  small  opening 
at  one  end,  called  a  blowpipe,  is  used  to  produce 
these  flames.  A  tube  with  a  flattened  top  is  put 
inside  the  burner  tube  to  produce  a  luminous  flame. 
The  tip  of  the  blowpipe  rests 
in  or  near  this  flame,  and  if 
air  is  gently  and  continuously 
blown  through  the  blowpipe, 
a  long,  slender  flame  is  pro- 
duced, called  a  blowpipe 
flame  (Fig.  44).  It  is  like 
the  Bunsen  flame  as  far  as 
its  oxidizing  and  reducing 

properties  are  concerned.  The  blowpipe  is  used 
in  the  laboratory  and  by  jewelers  and  mineral- 
ogists. On  a  large  scale  the  blowpipe  flame  is  used  to  reduce  or  oxidize 
ores  and  to  melt  refractory  substances  (see  Compound  Blowpipe) . 

The  Bunsen  flame  has  recently  been  utilized  in  producing  the  Wels- 
bach  light.  The  non-luminous  flame  heats  an  inverted  bag  or  "  man- 
tle "  of  oxides  of  rare  metals,  and  the  mantle  glows  with  an  intense 
light.  The  candle  power  varies  from  40  to  100.  This  form  of  burner 
is  widely  used  because  it  produces  a  brilliant  light. 

EXERCISES. 

1.  What  are  hydrocarbons  ?    ^  Where  are  they  found  ?     Name  sev- 
eral familiar  substances  containing  hydrocarbons. 

2.  Are  there  many  hydrocarbons  ?     Why  ? 

3.  What  is  an  homologous  series  of  hydrocarbons  ?      Name  four 
such  series. 


FIG.  43.  — The  oxi- 
dizing (A)  and  reduc- 
ing (Z?)  flames. 


FIG.  44.  —  Blowpipe 
flame,  showing  oxidiz- 
ing (A)  and  reducing 
(B)  parts. 


Exercises.  223 

4.  What  is  methane  ?    What   other  names  has   it  ?     Where  is  it 
found  ?     How  is  it  usually  prepared  ?     State  its  essential  properties. 
Why  is  it  a  dangerous  gas  ?     Illustrate  your  answer  by  an  equation. 

5.  What  other  name  has  the  methane  series  ?     Why  ?     Illustrate 
the  following  terms  by  the  paraffin  series :  (a)  substitution,  (£)  substi- 
tution product,  (6-)  saturated  hydrocarbon. 

6.  What  is  ethylene  ?     How  is  it  prepared  ?    Where  is  it  found  ? 
State  its  properties.     Give  the  equation  expressing  the  combustion  of 
ethylene. 

7.  Illustrate  the  following  terms  by  the  ethylene  series  :  (a)  unsatu- 
rated  hydrocarbon,  (b)  addition  product,  (c)  isomerism,  (d}  metamer- 
ism, (e)  isomer. 

8.  Review  the  subject  of  calcium  carbide  (see  Chapter  X). 

9.  What  is  acetylene  ?     How  is  it  formed  ?     How  is  it  prepared  ? 
Give  the  equation  for  the  reaction.     Summarize  the  properties  of  acety- 
lene. 

10.  Illustrate  the  following   terms   by  acetylene :    (a)  polymerism, 
(^)  polymer,  (V)  unsaturated    hydrocarbon. 

1 1 .  Describe  the  acetylene  (a)  flame,  (£)  burner,  and  (c)  generator. 
What  precautions  must  be  observed  in  using  acetylene  as  an  illuminant  ? 

12.  What   is    (a)   choke   damp,    (£)   black   damp,    (V)  marsh  gas, 
(//)  olefiant  gas  ? 

13.  What  is  the  formula  of  (a)  methane,  (£)  ethylene,  (c)  benzene  ? 
Why  is  C2H2  the  formula  of  acetylene4  ? 

14.  How  many  volumes  of  oxygen  are  needed  for  the  combustion  of 
one  volume  of  (a)  methane,  (£)  ethylene,  and  (c)  of  two  volumes  of 
acetylene  ?    What  volumes  of  what  products  are  formed  in  each  case  ? 
What  law  do  these  relations  illustrate  ? 

15.  What  is  petroleum  ?     Where  is  it  found  ?     Of  what  is  petroleum 
composed  ?     How  is  it  obtained  from  the  earth  ?     Describe  briefly  the 
refining  of  petroleum. 

1 6.  What  is  kerosene  ?     Describe  its  method  of  preparation.     Define 
and  illustrate  the  \ES\bfldskingpoint. 

17.  State  the  uses  of  (a)  gasoline,  (<£)  lubricating  oils,  (c)  vaseline, 
(tf)  paraffin  wax. 

1 8.  What  is  natural  gas  ?    Where  is  it  found  ?     Of  what  is  it  com- 
posed ?     For  what  is  it  used  ? 

19.  What  is  coal  gas  ?     Describe  briefly  its  manufacture. 

20.  What  is  coal  tar  ?    What  are  its  uses  ? 


224  Descriptive  Chemistry. 

21.  What  is  ammoniacal  liquor  ?     What  is  its  source  ?     How  is  it 
obtained  ?     For  what  is  it  used  ? 

22.  Review  (a)  coke,  and  (b}  gas  carbon  (see  Chapter  XIV). 

23.  What  is  water  gas  ?     Describe  briefly  its  manufacture.     What  is 
meant  by  "  enriching  "  water  gas  ?     What  is  producer  gas  ? 

24.  Give  the  equation  for  the  interaction  of  carbon  and  steam.    How 
many  volumes  of  steam  are  needed  to  produce  one  volume  of  each  of  the 
products  ? 

25.  What  is  illuminating  gas  ?      State  its  chief  properties.      What 
are  its  (a}  light-giving  constituents,  (b}  diluents,  (c)  impurities  ?    Upon 
what  does  its  luminosity  depend  ?     How  is  this  property  measured  and 
expressed  ?     Give  two  reasons  why  illuminating  gas  is  dangerous. 

26.  What  is  a  flame  ?     Illustrate  your  answer.     Describe  the  struc- 
ture of  a  candle  flame.    What  are  the  chief  gaseous  products  of  combus- 
tion ?     Why  do  lamps  sometimes  smoke  ?     What  affects  the  luminosity 
of  many  flames  ? 

27.  Describe  (a)  the  Bunsen  flame,  (fr)  the  Bunsen  burner.     Why 
is  the  Bunsen  flame  non-luminous  ?     Describe  and  explain  the  "  strik- 
ing back"  of  the  Bunsen  flame.     Describe  the  structure  of  the  Bunsen 
flame.     What  is  the  miner's  safety  lamp,  and  upon  what  principle  is  it 
constructed  ? 

28.  Review  oxidation  and  reduction. 

29.  What  is  (#)  an  oxidizing  flame  ?     Describe  a  blowpipe  and  its 
flame.     For  what  is  it  used  ? 

30.  Describe  the  Welsbach  light. 


PROBLEMS. 

1.  Calculate  the  percentage  composition  of  (a)  methane  (CH4), 
(£)  ethylene  (C2H4),  and  (c}  acetylene  (C2H2). 

2.  What  weight  of  oxygen  is  needed  for  the  complete  combustion 
of  4  gm.  of  ethylene ?     (Equation  is  C2H4  +  3<32  =  2  CO,  +  2  H2O.) 

3.  What  is  the  simplest  formula  of  a  compound  having  the  compo- 
sition H  =  7.69  and  C  =  92.3  ? 

4.  Calculate  the  molecular  formula  of  a  compound  having  the  vapor 
density  38.8  and  the  composition  C  =  92.3  and  H  =  7.69. 


CHAPTER    XVI. 
FLUORINE  -  BROMINE  —  IODINE. 

FLUORINE,  bromine,  and  iodine,  together  with  chlorine, 
are  often  grouped,  and  called  the  fialogenj.  They  resem- 
ble each  other  in  a  general  way,  aiuT  forni  analogous  com- 
pounds which  have  similar  properties,  differing  mainly  in 
degree. 

Halogen  means  "  a  sea-salt  producer."  It  is  applied  to  this  group 
of  elements  because  they  form  salts  which  resemble  sodium  chloride 
(common  sslt  or  sea  salt).  Chlorides,  bromides,  and  iodides  are  some- 
times called  haloid  salts  or  halides.  The  Greek  word  for  salt,  hals, 
suggested  these  terms. 

FLUORINE. 

Occurrence.  —  Fluorine  is  the  most  active  of  all  the  ele- 
ments, and  is  therefore  never  found  free  in  nature.  It 
occurs  abundantly  in  combination  with  calcium  as  fluor 
spar  or  calcium  fluoride  (CaF2).  Other  native  compounds 
are  cryolite  (Na3AlF6)  and  apatite  (CaF2.  3  Ca3(PO4)2). 
Minute  quantities  of  combined  fluorine  are  found  in  bones 
and  blood,  in  the  enamel  of  the  teeth,  and  in  sea  and  some 
mineral  waters. 

Fluorine  is  named  from  fluor  spar,  which  melts  easily  and  is  used  as 
a  flux  to  make  substances  flow  together  (hence  the  derivation  from  the 
Latin  fluo,  I  flow). 

The  Isolation  of  Fluorine  was  accomplished  in  1886  by 
Moissan,  though  many  unsuccessful  attempts  had  been 
previously  made.  He  decomposed  hydrofluoric  acid  by 

225 


226 


Descriptive  Chemistry. 


electricity  and  collected  the  liberated  fluorine.  The 
achievement  was  attended  with  tremendous  difficulties, 
owing  to  the  intense  activity  of  fluorine  and  its  corrosive 
properties. 

The  essential  parts  of  the  apparatus  used  by  Moissan  are  shown  in 
Figure  45.  The  U-tube,  made  of  an  alloy  of  platinum  and  iridium,  is 
provided  with  tightly  fitting  stoppers  of  fluor 
sr  (S.  S) .  Through  the  stoppers  pass  the  elec- 
trodes (E,  E)  of  platinum  iridium,  held  in  place 
by  screw  caps  (C,C).  Side  tubes  ( T,  T)  allow 
the  liberated  gases  (fluorine  and  hydrogen) 
to  be  drawn  off  separately  through  platinum 
delivery  tubes.  Perfectly  dry  hydrofluoric 
acid  is  put  into  the  U-tube  and  dry  acid 
potassium  fluoride  (HKF2)  is  added  to  enable 
the  solution  to  conduct  the  current  —  liquid 
hydrofluoric  acid  itself  being  a  non-conductor. 
The  U-tube  is  cooled  to  a  very  low  tempera- 
ture (—23°  to  —  50°  C.),  and  on  passing  a 
current  through  the  apparatus  fluorine  is 
evolved  at  the  positive  electrode  and  hydrogen 
at  the  other.  The  fluorine,  freed  from  hydro- 
fluoric acid  vapor,  was  collected  by  Moissan 
at  first  in  a  platinum  tube  with  thin  fluor  spar  plates  closing  each  end, 
so  that  he  could  look  inside  and  examine  the  gas.  Later  he  found  that 
pure  fluorine  can  be  collected  in  glass  tubes,  since  it  attacks  glass  only 
very  slowly. 

Properties.  —  Fluorine  has  a  sharp  odor  and  a  greenish 
yellow  color,  but  lighter  and  more  yellowish  than  chlorine. 
Its  density  is  1.265  (an"  =  0-  Subjected  to  pressure  and 
a  very  low  temperature,  it  condenses  to  a  pale  yellow  liquid, 
which  boils  at  —187°  C.  The  pure  gas  can  be  liquefied 
in  a  glass  vessel.  Chemically,  fluorine  is  intensely  active. 
Hydrogen,  bromine,  iodine,  sulphur,  phosphorus,  carbon, 
silicon,  and  boron  take  fire  in  it.  Oxygen,  nitrogen,  and 
argon  do  not  unite  with  it.  Most  metals  burn  in  it,  form- 


FlG.  45.  —  Moissan's  ap- 
paratus for  preparing  flu- 
orine. 


Fluorine  —  Bromine  —  Iodine.  227 

ing  fluorides.  Gold  and  platinum  are  not  attacked  by  it 
below  red  heat.  Copper  becomes  coated  with  copper  fluor- 
ide, which  protects  the  metal,  so  that  copper  vessels  may 
be  used  as  fluorine  generators.  Moissan  used  a  copper 
U-tube  to  prepare  large  volumes.  Water  is  decomposed 
by  it  at  ordinary  temperatures,  owing  to  the  intense  attrac- 
tion between  hydrogen  and  fluorine ;  hydrocarbons,  for  a 
similar  reason,  are  instantly  decomposed,  hydrofluoric  acid 
and  carbon  fluorides  being  the  products. 

The  exhaustive  work  of  Moissan  shows  that  fluorine,  though  more 
active  than  the  other  halogens,  is  similar  to  them,  and  should  be  regarded 
as  the  first  member  of  that  group. 

Hydrofluoric  Acid,  HF,  is  the  compound  of  fluorine 
corresponding  to  hydrochloric  acid.  It  is  prepared  by 
the  interaction  of  a  fluoride  and  concentrated  sulphuric 
acid.  Calcium  fluoride  is  usually  used,  and  the  experi- 
ment is  performed  in  a  lead  dish.  The  chemical  change 
is  represented  thus  — 

CaF2         +     H2SO4     =         2HF       +       CaSO4 

Calcium  Fluoride   Sulphuric  Acid   Hydrofluoric  Acid  Calcium  Sulphate 

Hydrofluoric  acid,  like  hydrochloric  acid,  is  a  colorless 
gas,  which  fumes  in  the  air  and  dissolves  in  water,  the 
solution  being  the  commercial  hydrofluoric  acid.  Both 
gas  and  liquid  are  dangerous  substances.  The  gas 
is  extremely  poisonous,  and  the  liquid,  if  dropped  on 
the  skin,  produces  terrible  sores.  Owing  to  its  corro- 
sive action  the  acid  is  preserved  and  sold  in  platinum, 
rubber,  or  wax  bottles.  The  acid  and  the  moist  gas  attack 
glass,  and  are  used  extensively  in  etching.  The  glass  is 
coated  with  wax,  and  the  design  to  be  etched  is  scratched 
through  the  wax.  The  glass  is  the*n  exposed  to  the  gas  or 
the  liquid,  which  attacks  the  exposed  places.  When  the 


228  Descriptive  Chemistry. 

wax  is  removed,  a  permanent  etching  like  the  design  is 
visible.  Glass  is  an  artificial  compound  of  silicon  —  a 
silicate.  The  corrosive  action  of  hydrofluoric  acid  upon 
glass  is  due  to  the  ease  with  which  the  acid  decomposes 
glass  and  forms  with  the  silicon  a  volatile  compound, 
called  silicon  tetrafluoride  (SiF4).  Since  silicon  dioxide 
(or  sand)  is  the  essential  constituent  of  the  mixture  from 
which  glass  is  made,  the  equation  for  etching  glass  may 
be  written  thus  — 

SiO2    +      4HF        =       SiF4       +  •  2  H2O 
Silicon         Hydrofluoric  Silicon 

Dioxide  Acid  Tetrafluoride 

Scales  on  thermometers  and  on  other  graduated  glass 
instruments  are  etched  with  hydrofluoric  acid. 

The  vapor  density  of  hydrofluoric  acid  gas  indicates  that  its  formula 
is  HF  at  high  temperature,  but  H2F2  at  lower  temperatures  (30°  C.). 

BROMINE. 

Occurrence.  —  Bromine  is  never  found  free  in  nature  on 
account  of  its  chemical  "activity.  Bromides  are  widely 
distributed,  especially  magnesium  bromide.  The  salt 
springs  of  Ohio,  West  Virginia,  Pennsylvania,  and  Michi- 
gan, and  the  salt  deposits  at  Stassfurt  in  Germany  furnish 
the  main  supply  of  the  element.  Sea  water,  Chili  salt- 
peter (NaNO3),  and  certain  seaweeds  contain  a  small 
quantity  of  combined  bromine. 

Preparation. —  Bromine  is  obtained  from  its  compounds 
by  treatment  with  chlorine,  or  with  sulphuric  acid  and 
manganese  dioxide.  In  the  laboratory,  bromine  is  pre- 
pared by  heating  potassium  bromide  with  manganese 
dioxide  and  sulphuric  a'cid  in  a  glass  vessel.  The  bromine 
is  easily  liberated  as  a  dense,  brown  vapor,  which  often 


Fluorine  —  Bromine  —  Iodine.  229 

condenses  to  a  liquid  and  runs  down  the  walls  of  the 
vessel.  The  chemical  change  is  represented  thus— - 

2  KBr  +  2  H2SO4+  MnO2  =  Br2  +  MnSO4+  K2SO4  +  2  H2O 

Potassium     Sulphuric      Manganese     Bro-      Manganese      Potassium        Water 
Bromide          Acid  Dioxide       mine       Sulphate         Sulphate 

Bromine  is  sometimes  prepared  by  treating  a  bromide  with 
manganese  dioxide  and  hydrochloric  acid. 

The  source  of  commercial  bromine  in  the  United  States  is  "  bittern  " 
—  a  concentrated  liquid  left  after  salt  is  crystallized  from  brine.  In  the 
continuous  process  the  hot  bittern  flows  down  a  large  tower  filled  with 
broken  brick  or  burned  clay  balls  ;  chlorine  gas  and  steam  forced  in  at 
the  bottom  meet  the  bittern  and  liberate  the  bromine,  which  passes  as  a 
vapor  out  of  the  top  into  a  condenser.  The  main  chemical  change  is 
represented  thus  — 

MgBr2  +        C12  -  Br2      -f  MgCl2 

Magnesium  Bromide       Chlorine  Bromine      Magnesium  Chloride 

In  the  periodic  process,  used  chiefly  in  the  United  States,  a  huge  stone 
still  is  charged  with  manganese  dioxide,  hot  bittern,  and  sulphuric  acid, 
and  heated  by  steam.  The  bromine  distills  into  a  condenser,  as  in  the 
other  process.  Sometimes  potassium  chlorate  is  used  as  the  oxidizing 
agent. 

Properties.VtJ^romineis  a  heavy,  reddish  brown  liquid 
at  the  ordinary  T^njgCIamel  Its  specific  gravity  is  about 
three.  It  is  a  volatile  liqu!?f,  boiling  at  about  59°  C.  The 
vapor,  which  is  given  off  freely,  has  a  disagreeable,  suffo- 
cating odor.  This  property  suggested  the  name  bromine 
(from  the  Greek  word  bromos,  a  stench).  It  is  poisonous, 
and  burns  the  flesh  frightfully.  Bromine  is  somewhat 
soluble  in  water.  The  solution,  called  bromine  water,  has 
a  brown  color,  and  when  cooled  deposits  a  crystalline 
hydrate  (Br2  .  10  H2O).  Many  other  properties  of  bromine 
are  similar  to  those  of  chlorine.  Thus,  it  combines  with 
metals  and  other  elements ;  it  also  bleaches. 


230  Descriptive  Chemistry. 

Compounds  of  Bromine  are  similar  to  those  of  chlorine.  Hydrobro- 
mic  acid  (HBr)  is  a  colorless,  pungent  gas,  which  fumes  in  the  air  and 
dissolves  freely  in  water,  forming  the  solution  usually  called  hydrobromic ' 
acid.  Its  other  properties  closely  resemble  those  of  hydrochloric  acid. 
Bromides  are  salts  of  hydrobromic  acid,  though  many  are  formed  by 
direct  combination  with  bromine.  Like  the  chlorides,  most  bromides 
dissolve  in  water.  Potassium  bromide  (KBr)  is  a  white  solid,  made  by 
decomposing  iron  bromide  with  potassium  carbonate.  It  is  used  exten- 
sively as  a  medicine  and  in  photography  (in  preparing  silver  bromide 
plates  and  films).  Bromides  of  sodium,  ammonium,  and  cadmium  have 
a  limited  use. 

Miscellaneous.  —  Bromine  itself  is  used  to  make  potassium  bromide 
and  other  compounds,  especially  a  class  of  coal  tar  dyes  used  to  color  pink 
string  and  to  make  red  ink.  Annually  over  500,000  pounds  of  bromine 
are  prepared  in  the  United  States,  while  Germany  exports  about  400,000 
pounds  of  bromine,  and  500,000  pounds  of  bromine  compounds. 

Balard  discovered  bromine  in  1826  in  the  mother  liquor  (or  bittern) 
from  brine.  Liebig  supposed  it  was  chloride  of  iodine,  and  thus  failed 
to  discover  it,  because,  as  he  said,  he  yielded  to  "  explanations  not 
founded  on  experiment." 

IODINE. 

Occurrence.  —  Free  iodine  is  never  found  in  nature,  but 
like  chlorine  and  bromine  it  is  combined  with  metals, 
especially  sodium,  potassium,  or  magnesium.  It  is  widely 
distributed,  though  the  quantity  in  any  one  place  is  small. 
Tobacco,  water  cress,  cod-liver  oil,  oysters,  and  sponges  con- 
tain minute  quantities.  Native  iodides  of  silver  and  of  mer- 
cury are  found.  The  ash  of  some  seaweeds  contains  from 
0.5  to  1.5  per  cent  of  its  weight  of  iodides  of  sodium  and 
potassium.  Sodium  iodate  occurs  in  the  deposits  of  salt- 
peter in  Chili,  and  is  now  the  main  source  of  the  element. 

Preparation.  —  Iodine  is  prepared  in  the  laboratory  by  a 
method  similar  to  that  used  for  bromine.  Potassium  iodide, 
manganese  dioxide,  and  sulphuric  acid  are  heated  in  a  glass 
vessel,  and  the  iodine  appears  as  a  violet  vapor,  which  con- 

A    . 


Fluorine — Bromine  —  Iodine.  231 

denses  on  the  upper  part  of  the  vessel  into  dark  grayish 
crystals. 

On  a  commercial  scale  iodine  is  prepared  from  the  ash  of  seaweeds 
and  from  the  mother  liquors  of  Chili  saltpeter,  (i)  Along  the  coasts 
of  France,  Scotland,  and  Norway  seaweed  is  collected  and  burned, 
usually  in  closed  vessels.  The  ash  is  called  kelp  or  varec.  The  solu- 
ble portions  are  removed  by  agitation  with  water.  The  'filtered  liquid 
is  further  purified,  and  from  the  final  mother  liquor  in  which  the  iodides 
are  dissolved,  the  iodine  is  extracted  by  heating  with  sulphuric  acid  and 
manganese  dioxide.  Sometimes  chlorine  is  used  to  extract  the  iodine. 
In  either  case  the  mother  liquor  and  its  added  ingredients  are  distilled 


FiG.  46.  —  Apparatus  for  purifying  iodine. 

gently  in  an  iron  pot  with  a  lead  cover,  which  is  connected  with  two 
rows  of  bottle-shaped  condensers  (Fig.  46).  The  iodine,  which  col- 
lects in  these  condensers,  is  purified  by  washing  and  resubliming. 
(2)  In  another  process  the  mother  liquor  from  the  Chili  saltpeter  is 
mixed  with  acid  sodium  sulphite  (HNaSO3),  and  the  precipitated  iodine 
is  collected  on  coarse  cloth,  washed,  dried,  and  then  resublimed,  as 
described  above. 

Courtois,  a  French  chemist,  discovered  iodine,  in  1812,  in  an  attempt 
to  prepare  potassium  nitrate  from  seaweed.  Davy  and  Gay-Lussac 
established  its  elementary  nature  and  discovered  many  of  its  properties. 
The  present  name  was  given  by  Davy. 

Properties.  —  Iodine  is  a  dark  grayish  crystalline  solid, 
resembling  graphite  in  luster.  It  crystallizes  in  plates 
which  have  the  specific  gravity  4.95.  It  is  volatile  at  the 


232  Descriptive  Chemistry. 

ordinary  temperature,  and  when  gently  heated  the  vapor 
which  is  formed  has  a  beautiful  violet  color.  This  color 
suggested  the  name  iodine  (from  the  Greek  word  iodes, 
violetlike).  The  vapor  is  nearly  nine  times  heavier  than 
air,  and  has  an  odor  resembling  dilute  chlorine,  though  less 
irritating.  When  the  vapor  is  heated,  its  color  changes 
from  violet  to  deep  blue,  and  the  density  decreases.  Ex- 
periment indicates  that  at  about  700°  C.  the  molecules  con- 
tain only  two  atoms,  and  as  the  temperature  rises  the 
molecules  dissociate,  until  at  a  very  high  temperature  the 
vapor  consists  entirely  of  atoms.  Iodine  stains  the  skin 
yellow,  and  turns  cold  starch  solution  blue.  The  presence 
of  a  minute  trace  of  iodine  may  be  thus  detected,  one  part 
of  iodine  in  over  400,000  parts  of  water  producing  the  blue 
color.  The  exact  nature  of  this  blue  compound  is  un- 
known. The  presence  of  starch  in  many  vegetable  sub- 
stances can  be  shown  by  this  delicate  test.  Iodine  dissolves 
slightly  in  water,  and  freely  in  alcohol,  chloroform,  carbon 
disulphide,  ether,  and  potassium  iodide  solution.  The 
chloroform  and  carbon  disulphide  solutions  are  violet,  but 
the  others  are  brown,  or  even  black.  The  chemical  proper- 
ties of  iodine  resemble  those  of  chlorine  and  bromine,  but 
it  is  less  active.  Bromine  and  chlorine  displace  iodine 
from  its  compounds,  chlorine  and  chlorine  water  being 
often  used  for  this  purpose.  It  combines  directly  with 
other  elements  and  replaces  some.  Phosphorus  bursts  into 
a  flame  when  mixed  with  iodine. 

Compounds  of  Iodine  resemble  the  corresponding  ones  of  chlorine 
and  bromine.  Hydriodic  acid  is  much  like  hydrobromic  and  hydro- 
chloric acid,  though  unlike  them  in  being  a  reducing  agent.  Iodides 
are  salts  of  hydriodic  acid,  and  like  many  salts  they  are  prepared  in 
various  ways.  In  general  behavior  they  are  similar  to  bromides  and 
chlorides.  Potassium  iodide  (KI)  is  made  and  used  like  potassium 
bromide.  lodates  and  periodates  are  known. 


Fluorine  —  Bromine-^- Iodine.  233 

Miscellaneous.  —  Iodine  dissolved  in  alcohol  or  in  potassium  iodide 
solution  is  used  as  an  application  for  the  skin  to  prevent  the  spread  of 
eruptions  or  to  reduce  swellings.  Iodine  is  used  to  make  medicinal 
preparations,  especially  iodoform  (CHI3),  which  is  used  as  a  dressing 
for  wounds.  Large  quantities  of  iodine  are  used  in  making  aniline 
dyes.  Potassium  iodide  is  made  in  large  quantities,  Germany  alone 
exporting  about  150  tons  of  it  annually.  Chili  annually  exports  over 
300  tons  and  Norway  over  160  tons  of  iodine  and  iodides. 


EXERCISES. 

1.  What  elements  constitute  the  halogen  group  ?     Why  are  they 
so  .called  ? 

2.  How  does  fluorine  occur  in  nature  ?     Describe  briefly  the  isola- 
tion of  fluorine.     When  was  it  first  performed?     Summarize  the  chief 
properties  of  fluorine. 

3.  How  is  hydrofluoric  acid  prepared?     Give  the  equation  for  the 
reaction.    What  are  its  characteristic  properties  ?     For  what  is  it  used  ? 

4.  How  is  glass  etched?     State  the  essential  changes. 

5.  What  is  the  formula  of  hydrofluoric  acid  ? 

6.  How  does  bromine  occur  in  nature  ?     What  are  the  sources  of 
commercial  bromine  ?     What  general  method  is  used  to  prepare  this 
element  ?     Describe  briefly  the  commercial  methods.     State  the  chief 
properties.     For  what  is  it  used  ?     How  does  this  element  differ  from 
all  others  previously  studied  ? 

7.  Name   several   compounds   of  bromine.      What   is   potassium 
bromide  ? 

8.  Give  a  brief  account   of  the   discovery  of  (a)  bromine   and 
(b)  iodine. 

9.  Discuss  the  occurrence  of  iodine  in  nature.     How  is  iodine  pre- 
pared (a}  in  the  laboratory  and  (b)  on  a  large  sqale  ?     Summarize  the 
properties  of  iodine.     Describe  the  test  for  iodine. 

10.    Name  several  compounds  of  iodine.     Describe  potassium  iodide, 
n.    Compare  hydrochloric,  hydrobromic,  and  hydriodic  acids. 

12.  What  is  the  symbol  of  (a}  fluorine,  (b}  chlorine,  (c)  bromine, 
(d)  iodine  ?     What  is  the  derivation  of  the  name  of  each  element  ? 

13.  Compare  the  physical  properties  of  fluorine,  chlorine,  bromine, 
and  iodine. 

14.  What  is  "  drug-store  iodine  "  ? 


234  Descriptive  Chemistry. 

PROBLEMS. 

1 .  What  is  the  percentage  composition  of  (a)  fluor  spar  (CaF2)  and 
(£)  cryolite  (Na3AlF6)? 

2.  How  much  (a)  calcium  sulphate  and  (£)  hydrofluoric  acid  are 
formed  by  heating  100  gm.  of  fluor  spar  with  sulphuric  acid  ? 

3.  Calculate  the  percentage  composition  of  (a)  potassium  bromide 
(KBr),  (£)  potassium  iodide  (KI),  (c}  silver  bromide  (AgBr),  and 
(</)  iodoform  (CHI3). 

4.  How  much  potassium  iodide  is  needed  to  prepare  63.5  gm.  of 
iodine  ? 

5.  How  much  potassium  bromide  is  needed  to  prepare  10  gm.  of 
bromine  ? 


CHAPTER   XVII. 
SULPHUR  AND  ITS  COMPOUNDS. 

SULPHUR  has  been  known  for  ages.  The  alchemists  re- 
garded it  as  one  of  the  primary  forms  of  matter.  The  ele- 
ment and  its  compounds  have  always  played  an  important 
part  in  the  development  of  many  industries. 

Occurrence  and  Formation.  —  Sulphur,  free  and  com- 
bined, is  abundant  and  widely  distributed.  Free  or  native 
sulphur  is  found  usually  in  volcanic  regions.  There  are 
also  beds  associated  with  gypsum  (calcium  sulphate).  It  is 
believed  that  such  deposits  were  formed  by  the  reduction  of 
the  gypsum  by  microorganisms  into  limestone  and  sulphur. 

Combined  sulphur  is  found  in  volcanic  gases,  in  sub- 
stances of  vegetable  and  animal  origin,  and  as  sulphides 
and  sulphates.  Several  important  metallic  ores  are  native 
sulphides,  e.g.  lead  sulphide  (PbS),  zinc  sulphide  (ZnS),  and 
those  of  mercury,  antimony,  and  copper.  Probably  some 
native  sulphur  has  been  formed  by  the  decomposition  of 
sulphides  by  heat.  The  most  abundant  sulphates  are 
varieties  of  calcium  sulphate  (CaSO4),  barium  sulphate 
(BaSO4),  and  magnesium  sulphate  (MgSO4).  Volcanic 
gases  often  contain  sulphur  dioxide  (SO2)  and  hydrogen 
sulphide  (H2S).  The  latter  is  also  found  in  the  water  of 
sulphur  springs.  Doubtless  some  of  the  sulphur  found  in 
volcanic  districts  has  been  produced  from  these  two  gases. 
Their  interaction  may  be  represented  thus  — 

SO2         +  2  H2S  =     3  S     +    2H2O 

Sulphur  Dioxide        Hydrogen  Sulphide         Sulphur          Water 

235 


236 


Descriptive  Chemistry. 


Sulphur  is  also  a  component  of  onions,  horse-radish,  mus- 
tard, garlic,  eggs,  some  petroleum  and  coal,  and  certain 
complex  compounds  of  the  body  —  such  as  bile  and  saliva. 
It  has  been  estimated  that  the  body  contains  about  125 
gm.  (0.27  Ib.)  of  combined  sulphur. 

Source.  —  Sicily  furnishes  most  of  the  sulphur  used  in 
the  world,  the  annual  output  being  about  500,000  tons. 
Owing  to  the  favorable  geographical  location,  rich  deposits, 
and  cheap  labor,  the  bulk  of  the  supply  will  continue  to 
come  from  this  island.  Some  sulphur  is  obtained  from 
Japan,  Italy,  Greece,  and  from  the  United  States,  especially 
in  Nevada,  Utah,  Idaho,  and  Louisiana. 

Some  of  the  sulphur  of  commerce  is  obtained  by  roasting  iron  pyrites, 
as  in  the  manufacture  of  sulphuric  acid.  Small  amounts  are  recovered 
from  the  calcium  sulphide  waste  of  the  Leblanc  soda  process  (see 
Sodium  Carbonate),  and  from  the  residues  of  the  iron  oxide  used  to 
purify  illuminating  gas.  . 


FIG.  47.  —  Kiln  for  extracting  sulphur  from  the  crude  ore.    The  calcarone  is  shown 
as  a  vertical  section  (right)  and  in  operation  (left). 

Extraction.  —  For  many  years  sulphur  has  been  ex- 
tracted from  the  impure  native  sulphur  in  Sicily  by  a 
primitive  process.  The  crude  sulphur  is  brought  to  the 
surface  by  laborers,  piled  loosely  in  a  heap,  and  covered 
with  powdered  or  burnt  ore  or  with  earth.  The  heap  is 
ignited  at  the  bottom,  and  the  heat  produced  by  the  com- 


Sulphur  and  its  Compounds. 


237 


bustion  of  some  of  the  sulphur  melts  the  rest,  which  runs 
out  at  the  bottom  (Fig.  47). 

This  method  is  being  discarded  in  the  more  prosperous  localities, 
because  it  is  wasteful  and  produces  intolerable  fumes.  Coal  instead  of 
sulphur  is  being  used  as  a  fuel,  and  extraction  by  hot  water  under 
pressure  is  coming  into  general  use.  In  some  cases  the  sulphur  is 
extracted  by  heating  the  crude  sulphur  with  a  hot  solution  of  calcium 
chloride. 

Purification.  —  Sulphur  obtained  from  its  ore  requires 
purification.  This  is  accomplished  by  the  apparatus  shown 
in  Figure  48.  The  crude  sulphur  is  melted  in  B,  and  flows 
into  the  iron  cylinder,  A.  Here  it  is  heated,  and  the  vapors 


FIG.  48.  — Apparatus  for  purifying  sulphur. 

pass  into  the  large  brick  chamber,  provided  with  a  tap,  C, 
from  which  the  liquid  sulphur  may  be  withdrawn.  If  the 
distillation  is  conducted  slowly,  the  sulphur  vapor  con- 
denses upon  the  cold  walls  of  the  chamber  as  a  fine 


238  Descriptive  Chemistry. 

powder,  called  flowers  of  sulphur,  just  as  water  vapor 
suddenly  cooled  below  o°  C.  turns  to  snow.  As  the 
operation  continues  the  walls  become  hot,  and  the  sulphur 
collects  on  the  floor  as  a  liquid  which  is  drawn  off  into 
wooden  molds.  This  is  roll  sulphur  or  brimstone. 

Properties,  —  Ordinary  sulphur  is  a  yellow,  brittle,  crys- 
talline solid.  It  is  insoluble  in  water,  but  most  varieties 
dissolve  in  carbon  disulphide,  and  to  some  extent  in  turpen- 
tine, chloroform,  and  benzene  (C6H6).  Sulphur  does  not 
conduct  heat.  The  warmth  of  the  hand  causes  it  to  crackle 
and  even  break  from  the  unequal  expansion. 

The  specific  gravity  of  the  solid  is  about  2.  The  specific  gravity  of 
the  vapor  varies  with  the  temperature.  At  the  lowest  temperature  at 
which  sulphur  can  be  vaporized,  the  molecule  contains  eight  atoms 
(S8),  while  at  900°  C.  and  higher  it  contains  two  atoms  (S2). 

Heated  to  1 14.5°  C.  sulphur  melts  to  a  thin,  amber-colored 
liquid.  As  the  temperature  is  raised,  the  liquid  darkens 
and  thickens,  until  at  about  230°  C.  it  is  black  and  too  thick 
to  be  poured  from  the  vessel.  Heated  still  higher,  the 
color  remains  black  but  the  mass  becomes  thin,  and  finally 
at  about  448°  C.  the  liquid  boils  and  turns  into  a  yellowish 
brown  vapor.  Sulphur  ignites  readily  and  burns  with  a 
pale  blue  flame,  forming  sulphur  dioxide  gas,  SO2;  if 
burned  in  oxygen,  a  little  sulphur  trioxide,  SO3,  is  also 
formed.  Finely  divided  sulphur  oxidizes  in  moist  air, 
forming  sulphuric  acid,  H2SO4.  It  also  combines  directly 
and  readily  with  hydrogen,  carbon,  chlorine,  and  other  ele- 
ments, especially  metals.  The  compounds  formed  are 
sulphides. 

The  reaction  between  sulphur  and  metals  is  often  attended  by  vivid 
combustion,  though  heat  is  necessary  to  start  the  chemical  action. 
When  a  mixture  of  flowers  of  sulphur  and  powdered  iron  is  heated,  the 
mass  begins  to  glow  and  soon  becomes  red-hot,  the  glow  often  spread- 


Sulphur  and  its  Compounds.  239 

ing  through  the  mass  after  removal  from  the  flame.  The  product  is 
iron  sulphide,  and  the  change  is  represented  thus  — 

Fe       +  S  FeS 

Iron  Sulphur        Iron  Sulphide 

Heated  copper  glows  when  dropped  into  melted  sulphur,  while  zinc 
dust  and  flowers  of  sulphur  combine  with  almost  explosive  violence. 

Different  Forms  of  Sulphur.  —  Sulphur  exists  in  at  least 
three  different  forms, — two  crystallized  and  one  amorphous. 
These  modifications  differ  in  specific  gravity,  solubility, 
and  other  properties.  The  crystallized  forms  belong  to 
the  orthorhombic  and  monoclinic  systems  (see  Appendix, 
§  3).  According  to  some  authorities  these  different  forms 
are  allotropic  modifications  of  sul- 
phur. Orthorhombic  sulphur  is  the 
form  deposited  by  crystallization 
from  a  solution  of  carbon  disulphide 
(Fig.  49).  Crystallized  native  sul- 
phur is  orthorhombic.  The  mono- 
clinic  crystals  are  deposited  from 
molten  sulphur.  By  melting  sul- 
phur in  a  crucible  and  pouring  off  FIG.  49.— Orthorhombic 
the  excess  of  liquid  as  soon  as  crys-  sulphur, 

tals  shoot  out  from  the  walls  near  the  surface,  the  interior  of 
the  crucible  when  cold  will  be  found  to  be  full  of  long,  dark 
yellow,  shining  needles.  They  are  monoclinic  crystals  of 
sulphur.  After  a  few  days  they  become  dull  and  yellow, 
and  crumble  into  minute  crystals  of  the  orthorhombic  form. 

Amorphous  sulphur  is  formed  by  pouring  boiling  sul- 
phur into  water.  It  is  a  tough,  plastic,  rubberlike,  amber- 
colored  mass,  insoluble  in  carbon  disulphide.  It  is  entirely 
different  in  color  and  texture  from  the  crystallized  varieties. 
In  a  short  time  it  becomes  hard,  brittle,  and  yellow,  like 
ordinary  sulphur. 


240  Descriptive  Chemistry. 

Other  varieties  of  amorphous  sulphur  are  known.  They  are  white 
or  whitish  powders.  One  is  made  by  boiling  flowers  of  sulphur  with 
milk  of  lime  and  adding  hydrochloric  acid  to  the  decanted  liquid.  A 
fine  sulphur  powder  is  precipitated,  which  gives  the  liquid  the  appear- 
ance of  milk,  hence  the  name  often  applied  to  it,  "milk  of  sulphur." 

Uses.  —  Sulphur  is  used  in  making  sulphuric  acid  and 
other  sulphur  compounds,  gunpowder,  fireworks,  matches, 
in  vulcanizing  rubber,  as  a  medicine  and  a  constituent  of 
some  ointments,  and  as  a  germicide  for  Phylloxera  —  an 
insect  which  destroys  grapevines. 

Compounds  of  Sulphur.  — The  important  compounds  of 
sulphur  are  hydrogen  and  other  sulphides,  sulphur  dioxide 
and  trioxide,  the  sulphites,  sulphuric  acid  and  the  sulphates, 
and  carbon  disulphide. 

Hydrogen  Sulphide,  H2S,  is  a  gaseous  compound  of 
sulphur  and  hydrogen,  and  is  often  called  sulphuretted 
hydrogen.  It  occurs  in  some  volcanic  gases,  and  in  the 
waters  of  sulphur  springs.  It  is  often  found  in  the  air, 
especially  near  sewers  and  cesspools,  since  it  is  one  prod- 
uct of  the  decay  of  organic  substances  containing  sul- 
phur. It  is  one  of  the  impurities  of  illuminating  gas,  being 
formed  by  the  union  of  the  sulphur  and  hydrogen  of  the 
coal. 

The  gas  is  prepared  in  the  laboratory  by  the  interaction 
of  dilute  acids  and  metallic  sulphides,  usually  hydrochloric 
acid  and  ferrous  sulphide.  When  the  acid  is  poured  upon 
fragments  of  the  sulphide,  the  gas  is  rapidly  evolved  with- 
out applying  heat,  and  may  be  collected  over  water.  The 
equation  for  the  chemical  change  is  — 

FeS       +       2HC1       =       H2S       +        FeCl2 

Iron  Hydrochloric          Hydrogen  Iron 

Sulphide  Acid  Sulphide  Chloride 


Sulphur  and  its  Compounds.  241 

Hydrogen  sulphide  gas  is  colorless  and  has  the  odor  of 
rotten  eggs.  It  is  poisonous.  A  little,  if  breathed,  produces 
headache  and  nausea,  and  a  large  quantity  renders  one  un- 
conscious. This  gas  is  inflammable  and  burns  with  a  bluish 
flame,  forming  water  and  sulphur  dioxide,  thus  — 

2H2S       +       3O2       =       2SO2       +       2H2O 
Hydrogen  Sulphide       Oxygen          Sulphur  Dioxide          Water 

If  the  supply  of  air  is  insufficient,  combustion  is  incom- 
plete and  sulphur  is  also  formed.  It  is  a  powerful  reduc- 
ing agent,  and  is  often  used  as  such  in  chemical  analysis. 
Even  sulphuric  acid  is  reduced  by  it,  thus  — 

H2S04     +      H2S      =      S02      +      S      4-      2  H2O 

Sulphuric  Hydrogen  Sulphur          Sulphur  Water 

Acid  Sulphide  Dioxide 

Hydrogen  sulphide  is  soluble  in  water,  one  volume  of  water 
dissolving  about  three  volumes  of  the  gas  at  the  ordinary 
temperature.  The  solution  is  called  hydrogen  sulphide 
water,  and  is  often  used  instead  of  the  gas.  The  solution 
reddens  litmus  and  decomposes  slowly,  sulphur  being  de- 
posited. 

A  liter  of  dry  hydrogen  sulphide  gas,  under  standard  conditions, 
weighs  1.542  gm.  When  metals  are  heated  in  dry  hydrogen  sul- 
phide, metallic  sulphides  are  formed  and  the  volume  of  hydrogen  liber- 
ated is  the  same  as  the  original  volume  of  gas.  Since  the  hydrogen 
molecule  is  H2,  there  must  be  two  atoms  of  hydrogen  in  the  hydrogen 
sulphide  molecule.  Its  vapor  density  is  17.15,  hence  the  molecular 
weight  is  34.3.  Subtracting  2  for  H2,  the  remainder  32.3  agrees  well 
with  the  atomic  weight  of  sulphur.  Hence,  there  can  be  only  one 
atom  of  sulphur  in  hydrogen  sulphide,  and  formula  must  be  H2S. 

Sulphides  may  be  regarded  as  salts  of  the  weak  acid, 
hydrogen  sulphide,  though  they  are  not  always  prepared 
directly  from  hydrogen  sulphide.  They  may  be  produced 


242  Descriptive  Chemistry. 

by  the  direct  union  of  sulphur  and  metals,  as  in  the  case 
of  iron  and  copper  sulphides  previously  mentioned,  or  by 
exposing  the  metal  to  the  moist  gas.  A  more  common 
way  is  to  precipitate  them  by  passing  the  gas  into  solutions 
of  metallic  compounds,  or,  sometimes,  by  adding. hydrogen 
sulphide  water.  Copper,  tin,  lead,  and  silver  are  rapidly 
tarnished  by  the  gas.  Silverware,  on  this  account,  turns 
brown  or  black,  especially  in  houses  heated  by  coal  and 
lighted  by  coal  gas,  because  hydrogen  sulphide  is  one 
product  of  the  combustion  of  coal  and  gas.  The  brown 
silver  sulphide  also  coats  silver  spoons  which  are  put  into 
mustard  or  eggs.  Lead  compounds  are  blackened  by  this 
gas,  owing  to  the  formation  of  lead  sulphide,  thus  — 

PbO       +        H2S  PbS       +   H20 

Lead  Oxide     Hydrogen  Sulphide     Lead  Sulphide     Water 

For  this  reason  houses  painted  with  "  white  lead  "  paint 
often  become  dark,  and,  similarly,  oil  paintings  are  dis- 
colored. The  blackening  of  a  solution  of  a  lead  com- 
pound is  the  customary  test  for  hydrogen  sulphide. 

Many  sulphides  have  a  brilliant  color.  Arsenious  sulphide  is  pale 
yellow,  cadmium  sulphide  is  golden  yellow,  manganese  sulphide  is  flesh 
colored,  zinc  sulphide  is  white,  antimony  sulphide  is  orange  red.  They 
vary  in  solubility.  The  sulphides  of  lead,  silver,  copper,  and  some 
other  metals  are  insoluble  in  dilute  hydrochloric  acid.  The  sulphides 
of  iron,  zinc,  and  some  other  metals  are  decomposed  by  dilute  hydro- 
chloric acid,  but  are  precipitated  if  ammonium  hydroxide  is  present. 
Sulphides  of  certain  metals  dissolve  in  water.  Hence  by  precipitating 
metals  under  different  conditions,  groups  of  metals  may  be  separated 
and  subjected  to  further  tests.  The  color  often  affords  a  ready  means 
of  detecting  each  sulphide.  Hydrogen  sulphide  is  thus  a  serviceable 
reagent  in  the  branch  of  chemistry  called  Qualitative  Analysis. 

Sulphur  Dioxide,  SO2,  is  the  common  compound  of  sul- 
phur and  oxygen.  It  occurs  in  the  gases  of  volcanoes, 


Sulphur  and  its  Compounds.  243 

and  to  a  slight  extent  in  the  atmosphere,  since  it  is  the 
usual  product  of  the  combustion  of  sulphur  and  sulphur 
compounds. 

When  sulphur  burns  in  air  (or  oxygen),  sulphur  dioxide 
is  formed,  thus  — 

S      +      O2  SO2 

Sulphur         Oxygen  Sulphur  Dioxide 

It  is  also  formed  by  roasting  iron  disulphide  (iron  pyrites) 
in  the  air,  thus  — 

2  FeS2       +      1 1  O     =        4  SO2       +       Fe2O3 

Iron  Disulphide         Oxygen         Sulphur  Dioxide         Iron  Oxide 

The  above  reaction  is  utilized  on  a  large  scale  in  the  com- 
mercial manufacture  of  sulphuric  acid. 

Sulphur  and  carbon  reduce  sulphuric  acid  to  sulphur  dioxide,  thus  — 

S        +        2  H2SO4  3  SO2       +.       2  H2O 

Sulphur  Sulphuric  Acid  Sulphur  Dioxide 

C     4-     2H2SO4     =      2SO2      +      CO2      +      2H2O 
Carbon  Carbon  Dioxide 

Two  methods  of  preparation  are  used  in  the  laboratory. 

(1)  If  copper  and  concentrated  sulphuric  acid  are  heated, 
a  series  of  complex  changes  results  finally  in  the  evolution 
of  sulphur  dioxide.     The  equation  is  usually  written  — 

Cu    +    2H2SO4    =         SO2         +       CuSO4  +  2H2O 
Copper     Sulphuric  Acid     Sulphur  Dioxide     Copper  Sulphate 

(2)  Dilute  sulphuric  (or  hydrochloric)  acid  dropped  upon 
a  sulphite  yields  sulphur  dioxide,  thus  — 

Na2SO3   +    H2SO4     =      SO2     +     Na2SO4  +  H2O 

Sodium  .          Sulphuric  Sulphur  Sodium 

Sulphite  Acid  Dioxide  Sulphate 


244  Descriptive  Chemistry. 

This  method  is  convenient  for  liberating  a  steady  current 
of  the  gas. 

Sulphur  dioxide  gas  has  no  color.  Its  odor  is  suffocating, 
being  the  well-known  odor  associated  with  burning  sulphur 
matches.  It  will  not  burn  in  the  air,  nor  will  it  support 
ordinary  combustion.  A  burning  taper  or  stick  of  wood 
is  instantly  extinguished  by  it,  but  finely  divided  metals, 
iron  for  example,  burn  in  it.  It  is  a  heavy  gas,  the  high 
density  (2.2)  allowing  it  to  be  readily  collected  by  down- 
ward displacement.  Low  temperature  and  pressure  change 
it  into  a  transparent,  colorless  liquid,  which  boils  at  —  8°  C. 
and  freezes  at  —  76°  C.  into  a  transparent,  icelike  solid. 
It  is  very  soluble  in  water.  At  the  ordinary  temperature 
one  volume  of  water  dissolves  about  forty  volumes  of  gas, 
but  loses  it  all  by  boiling.  This  solution  is  sour  and  red- 
dens blue  litmus,  and  contains  sulphurous  acid.  Moist 
sulphur  dioxide  bleaches  vegetable  coloring  matters.  A 
red  or  a  purple  flower  loses  color  in  it.  Silk,  hair,  straw, 
wool,  and  other  delicate  substances,  which  would  be  injured 
by  chlorine,  are  whitened  by  sulphur  dioxide.  In  some 
cases  the  color  returns  when  the  bleached  article  is  exposed 
to  the  air  for  some  time,  and  usually  such  bleached  objects 
become  yellow  with  age.  The  coloring  matter  is  not  wholly 
destroyed,  but  probably  unites  with  the  sulphur  dioxide  to 
form  a  colorless  compound,  which  slowly  decomposes. 

Immense  quantities  of  sulphur  dioxide  are  used  in  the 
manufacture  of  sulphuric  acid.  The  gas  is  also  used  to 
preserve  meat  and  wines,  to  fumigate  clothing  and  houses, 
in  paper  making,  in  tanning,  in  refining  sugar,  and  in 
making  acid  sodium  sulphite.  Liquid  sulphur  dioxide  is  used 
in  extracting  glue  and  gelatine,  and  in  various  metallurgical 
processes.  It  absorbs  heat  during  evaporation,  and  is  used 
in  some  ice  machines. 


Sulphur  and  its  Compounds.  245 

A  liter  of  sulphur  dioxide  under  standard  conditions  weighs  2.868  gm. 

The  Composition  of  Sulphur  Dioxide  is  based  on  the  following: 
The  gas  formed  by  burning  sulphur  in  a  measured  volume  of  oxygen 
has  the  same  volume  as  the  oxygen  itself.  Hence  there  are  as  many 
molecules  of  sulphur  dioxide  as  there  were  of  oxygen ;  that  is,  one 
molecule  of  sulphur  dioxide  contains  one  molecule  (or  two  atoms)  of 
oxygen.  A  molecule  of  oxygen  weighs  32.  But  the  molecular  weight 
of  sulphur  dioxide  found  from  its  vapor  density  is  about  64.  Subtract- 
ing 32  (i.e.  2  x  1 6)  from  this,  there  remains  about  32  for  sulphur.  The 
atomic  weight  of  sulphur  is  32.07,  hence  sulphur  dioxide  contains  only 
one  atom  of  sulphur,  and  its  composition  is  expressed  by  the  formula 
S02. 

Sulphurous  Acid  and  Sulphites.  —  Sulphurous  acid  is  formed  when 
sulphur  dioxide  dissolves  in  water.  Sulphur  dioxide  is,  therefore,  sul- 
phurous anhydride.  The  simplest  equation  expressing  this  fact  is  — 

SO,  +          H2O  H2S03 

Sulphur  Dioxide  Water  Sulphurous  Acid 

The  acid  has  never  been  obtained  free,  resembling  carbonic  acid  in  this 
respect.  It  is  unstable,  and  gradually  forms  sulphuric  acid  by  combin- 
ing with  oxygen  from  the  air.  The  acid  is  dibasic,  and  forms  two 
classes  of  salts,  the  sulphites.  They  are  reducing  agents,  and  yield 
sulphur  dioxide  when  treated  with  acids.  Acid  sodium  sulphite 
(HNaSO3),  often  called  bisulphite  of  soda,  is  the  antichlor  used  to 
remove  the  excess  of  chlorine  from  bleached  cotton  cloth.  It  is  also 
used  in  brewing,  tanning,  and  in  making  starch,  sugar,  and  paper. 
Acid  calcium  sulphite  (CaH2(SO3)2),  prepared  by  passing  sulphur 
dioxide  into  milk  of  lime,  is  used  in  paper  making. 

Sulphur  Trioxide,  SO3,  is  formed  by  the  direct  union  of 
sulphur  dioxide  and  oxygen,  a  little  being  produced  when 
sulphur  burns  in  air  or  in  oxygen.  The  action  is  slow,  but 
may  be  hastened  by  passing  a  mixture  of  sulphur  dioxide 
and  oxygen  (or  air)  over  hot  platinum,  or  over  asbestos 
coated  with  platinum.  Other  substances  also  hasten  the 
change.  It  is  a  white,  crystalline  solid,  which  melts  at 
15°  C.  and  boils  at  46°  C.  Another  form,  silklike  in  luster 
and  appearance,  is  known.  When  exposed  to  moist  air  it 


246  Descriptive  Chemistry. 

fumes  strongly,  forming  sulphuric  acid  ;  and  when  dropped 
into  water  it  dissolves  with  a  hissing  sound  and  evolution 
of  heat,  thus  — 

SO3  +        H2O  H2SO4 

Sulphur  Trioxide  Water  Sulphuric  Acid 

The  vapor  density  of  sulphur  trioxide  shows  that  its  molecular  weight 
is  about  80.  Hence  the  formula  (SO3)  harmonizes  with  the  fact  that 
two  volumes  of  sulphur  trioxide  decompose  by  heat  into  two  volumes  of 
sulphur  dioxide  and  one  volume  of  oxygen. 

Sulphuric  Acid,  H2SO4,  is  found  in  the  waters  of  a  few 
rivers  and  mineral  springs.  It  is  manufactured  in  enormous 
quantities  and  used  for  many  purposes. 

Sulphuric  acid  was  doubtless  known  to  the  Arabian  alchemists  living 
in  the  tenth  century.  It  was  definitely  mentioned  by  Basil  Valentine 
in  the  fifteenth  century,  who  describes  its  preparation  by  heating  a  mix- 
ture of  iron  sulphate  (green  vitriol)  and  sand.  The  product,  an  oily 
liquid,  was  called  oil  of  vitriol,  a  name  now  often  used.  About  1740, 
the  method  of  burning  sulphur  and  oxidizing  the  product  was  introduced 
into  England. 

The  Manufacture  of  Sulphuric  Acid,  as  usually  con- 
ducted, is  based  upon  the  fact  that  the  oxidation  of  sulphur 
dioxide  in  the  presence  of  water  forms  sulphuric  acid. 
The  apparent  equation  for  the  chemical  change  is  — 

S02         4-0          4-     H20     =         H2S04 

Sulphur  Dioxide      Oxygen  Water  Sulphuric  Acid 

The  oxidation  is  accomplished  in  the  older  and  more  com- 
mon method  by  oxides  of  nitrogen.  T)  * 

The  general  operation  consists  in  passing  sulphur  diox- 
ide, air,  steam,  and  oxides  of  nitrogen  into  large  lead  cham- 
bers. The  oxides  of  nitrogen  in  the  presence  of  steam 
change  the  sulphur  dioxide  into  sulphuric  acid,  which  col- 
lects on  the  walls  and  floors  of  the  lead  chambers.  The 
oxides  of  nitrogen  which  lose  part  of  their  oxygen  by  this 


Sulphur  and  its  Compounds. 


247 


change  are  themselves  reoxidized  by  the  air  into  higher 
oxides,  thus  being  fitted  to  oxidize  more  sulphur  dioxide. 
The  oxides  of  nitrogen  act  as  carriers  of  oxygen,  continu- 


ously  giving  oxygen  to  sulphur  dioxide  and  taking  it  from 
the  air.  Theoretically,  a  small  quantity  of  the  oxides  of 
nitrogen  will  change  an  infinite  quantity  of  sulphur  dioxide 


248  Descriptive  Chemistry. 

into  sulphuric  acid,  but  in  practice  losses  occur  and  oxides 
of  nitrogen  must  be  supplied.  The  main  parts  of  a  sul- 
phuric acid  plant,  together  with  the  courses  taken  by  the 
gases,  are  shown  in  Figure  50. 

Careful  study  shows  that  the  chemical  changes  involved  in  this  pro- 
cess of  manufacturing  sulphuric  acid  are  complex  and  variable.  Ac- 
cording to  a  reliable  authority,  the  main  continuous  reactions  may  be 
represented  thus  — 

2HNO3     +     2SO2     +     H20     =     2H2SO4     +     N2O3 

2SO2         -t-     N2O3      -f     O2        +     H2O  =  2  SO2(OH)(NO2) 

Nitrosy  1-sul  phuric 

Acid 
2SO2(OH)(NO2)   +   H2O=   2H,SO4       +     N,O3 

or,    2SO2(OH)(NO2)   -f   SO2  +  O  +  2  H2O  =  3  H2SO4  +   N,O3 

The  nitrogen  trioxide  (N2O3)  is  the  essential  factor,  though  probably 
the  change  is  really  due  to  a  mixture  of  nitric  oxide  (NO)  and  nitrogen 
peroxide  (NO2) .  Under  some  conditions,  nitric  oxide  plays  a  prominent 
part.  It  may  be  said  in  general  that  the  ease  with  which  the  oxides  of 
nitrogen  pass  into  each  other  makes  it  highly  probable  that  they  are 
carriers  of  oxygen  from  the  air  to  the  sulphur  dioxide. 

A  Sulphuric  Acid  Plant  consists  of  three  main  parts — (a)  the  furnace 
for  producing  sulphur  dioxide,  (£)  the  lead  chambers  together  with  the 
Glover  and  Gay-Lussac  towers  for  changing  the  sulphur  dioxide  into 
sulphuric  acid,  and  (V)  the  concentrating  apparatus.  The  manufacture 
is  conducted  somewhat  as  follows:  (i)  Sulphur  or  iron  disulphide 
(FeS2)  is  burned  in  a  furnace  constructed  so  that  enough  air  passes 
over  the  burning  mass  to  change  the  sulphur  into  sulphur  dioxide,  and 
to  furnish  the  proper  amount  of  oxygen  for  later  changes.  In  some 
works  the  furnace  is  provided  with  "  niter  pots  "  containing  a  mixture 
of  sodium  nitrate  and  sulphuric  acid ;  the  nitric  acid  vapors  which  are 
formed  are  one  source  of  the  oxides  of  nitrogen.  (2)  The  mixture  of 
sulphur  dioxide,  oxides  of  nitrogen,  and  air  passes  from  the  furnace  into 
the  bottom  of  the  Glover  tower.  This  is  a  tall  tower  filled  with  small 
stones  over  which  flow  two  streams  of  sulphuric  acid,  one  dilute  and  the 
other  containing  oxides  of  nitrogen  (obtained  from  the  Gay-Lussac 


Sulphur  and  its  Compounds.  249 

tower).  These  acids  not  only  cool  the  ascending  gases,  but  are  them- 
selves deprived  of  water  and  oxides  of  nitrogen.  Hence,  concen- 
trated acid  flows  out  of  the  bottom  of  the  Glover  tower,  while  from  the 
top  sulphur  dioxide,  oxides  of  nitrogen,  steam,  and  air  pass  on  into  the 
first  lead  chamber.  Here  nitric  acid  is  often  introduced,  as  well  as 
steam.  The  main  chemical  changes  occur  in  this  and  in  the  second 
chamber.  A  third  chamber  serves  mainly  to  cool  and  dry  the  gases. 
These  chambers  are  huge  boxes  often  having  a  total  capacity  of  150,000 
cubic  feet ;  the  walls  and  floors  are  of  sheet  lead  supported  on  a  wooden 
framework,  lead  being  a  metal  which  is  only  slightly  attacked  by  the 
chamber  acid.  The  remaining  gases  pass  on  into  the  bottom  of  the 
Gay-Lussac  tower.  This  tower  is  filled  with  coke  over  which  flows 
concentrated  sulphuric  acid  (from  the  Glover  tower),  which  absorbs  the 
unused  oxides  of  nitrogen.  These  oxides  are  liberated  again  in  the 
Glover  tower,  hence  there  is  little  loss.  At  the  end  of  the  plant  is  a 
tall  chimney,  which  serves  as  an  exit  for  unused  gases  (such  as  nitro- 
gen) and  also  creates  a  draft  strong  enough  to  carry  the  gases  through 
the  chambers  and  tower.  (3)  The  acid  which  is  produced  in  the  cham- 
bers and  drawn  off.  from  them  at  intervals  contains  about  67  per  cent 
of  the  compound  H2SO4.  Ordinary  commercial  sulphuric  acid  which 
contains  about  96  to  98  per  cent  is  prepared  from  the  chamber  acid  by 
evaporation,  first  in  lead  pans  and  finally  in  a  platinum  or  an  iron 
vessel. 

Another  method  of  manufacturing  sulphuric  acid  has 
recently  been  perfected,  called  the  contact  method.  Sul- 
phur dioxide  and  air,  carefully  purified  and  properly  cooled, 
are  led  through  pipes  containing  plates  covered  with  a 
contact  mixture,  which  is  chiefly  finely  divided  platinum. 
The  sulphur  dioxide  is  oxidized  to  sulphur  trioxide,  thus  — 

SO2  +        O         =  SO3 

Sulphur  Dioxide  Oxygen  Sulphur  Trioxide 

The  sulphur  trioxide  is  conducted  into  dilute  sulphuric  acid 
or  water,  thus  producing  a  pure  acid  of  any  desired  strength. 
The  process  is  continuous  if  the  gases  from  the  pyrites 
burners  are  completely  freed  from  arsenic  compounds,  sul- 
phur dust,  and  other  impurities. 


250  Descriptive  Chemistry. 

In  the  above  process  the  platinum  is  not  changed,  nor  does  it  cause 
the  sulphur  dioxide  to  unite  with  the  oxygen.  It  facilitates  the  chemi- 
cal action  between  the  gases  somewhat  as  oil  assists  the  movement  of 
machinery.  This  kind  of  chemical  action  is  called  catalysis  or  cata- 
lytic action.  The  substance  which  hastens  or  retards  a  chemical 
reaction,  but  appears  unchanged  at  the  end  of  the  process  is  called  a 
catalyzer.  In  many  cases  of  catalytic  action  it  has  been  found  that  the 
catalyzer  probably  participates  in  the  chemical  action,  though  its  exact 
share  is  not  always  clearly  understood. 

Properties  of  Sulphuric  Acid.  —  Sulphuric  acid  is  an 
oily  liquid,  colorless  when  pure,  but  usually  brown  from 
the  presence  of  charred  organic  matter,  such  as  dust  and 
straw.  The  commercial  acid  has  the  specific  gravity  1.83. 
When  sulphuric  acid  is  mixed  with  water,  considerable 
heat  is  evolved.  The  acid  should  always  be  poured  into 
the  water,  otherwise  the  intense  heat  may  crack  the  vessel 
or  spatter  the  hot  acid.  The  volume  of  dilute  acid  pro- 
duced is  smaller  than  the  sum  of  the  volumes  of  water  and 
concentrated  acid.  The  tendency  to  absorb  water  is  shown 
in  many  ways.  The  concentrated  acid  absorbs  moisture 
from  the  air  and  from  gases  passed  through  it.  It  is  often 
used  in  the  laboratory  to  dry  gases,  since  it  is  not  volatile 
at  the  ordinary  temperature.  Wood,  paper,  sugar,  starch, 
cotton  cloth,  and  many  organic  substances  are  blackened  by 
sulphuric  acid.  Such  compounds  contain  hydrogen  and 
oxygen  in  the  proportion  to  form  water ;  these  two  ele- 
ments are  abstracted  and  carbon  alone  remains.  Similarly, 
sulphuric  acid  withdraws  water  from  the  flesh,  making 
painful  wounds. 

Sulphuric  acid  is  reduced  by  hydrogen  sulphide,  hydrobromic  and 
hydriodic  acids,  carbon,  and  sulphur ;  it  combines  with  ammonia  to  form 
ammonium  sulphate  (NH4)2SO4;  and  is  decomposed  by  all  metals  ex- 
cept platinum  and  gold,  liberating  hydrogen,  sulphur  dioxide,  or  hydro- 
gen sulphide. 


Sulphur  and  its  Compounds.  251 

Uses  of  Sulphuric  Acid.  —  Sulphuric  acid  is  one  of  the 
most  important  substances.  Directly  or  indirectly  it  is 
used  in  hundreds  of  industries  upon  which  the  comfort, 
prosperity,  and  progress  of  mankind  depend.  It  is  used 
in  the  manufacture  of  all  other  mineral  acids  and  many 
organic  acids.  It  is  essential  in  one  process  for  the  manu- 
facture of  sodium  carbonate,  from  which  in  turn  are  made 
soap  and  glass.  Enormous  quantities  are  consumed  in 
making  artificial  fertilizers,  alum,  nitroglycerine,  glucose, 
phosphorus,  dyestuffs,  and  in  various  parts  of  such  funda- 
mental industries  as  dyeing,  bleaching,  electroplating, 
refining,  and  metallurgy. 

Sulphates.  — JSulphuric  acid  is  dibasic  and  forms  two 
classes  of  salts,  —  the  normal  sulphates,  such  as  Na2SO4, 
and  the  acid  sulphates,  such  as  HNaSO4.  The  normal 
sulphates  are  stable  salts ;  the  acid  salts  lose  water  when 
heated.  Most  sulphates  are  soluble  in  water,  only  the  sul- 
phates of  barium,  strontium,  and  lead  being  insoluble, 
while  calcium  sulphate  is  slightly  soluble.  Important 
sulphates  are  calcium  sulphate  (gypsum  CaSO4.2  H2O), 
barium  sulphate  (heavy  spar,  BaSO4),  zinc  sulphate  (white 
vitriol,  ZnSO4),  copper  sulphate  (blue  vitriol  or  blue  stone, 
CuSO4),  iron  sulphate  (green  vitriol,  copperas,  ferrous  sul- 
phate, FeSO4),  sodium  sulphate  (Glauber's  salt,  Na2SO4), 
and  magnesium  sulphate  (Epsom  salts,  MgSO4).  Sul- 
phates are  widely  used  in  medicine  and  in  many  industries. 

The  test  for  sulphuric  acid  or  a  soluble  sulphate  is  the  formation 
of  the  white,  insoluble  barium  sulphate  upon  the  addition  of  barium 
chloride  solution.  An  insoluble  sulphate  fused  on  charcoal  is  reduced 
to  a  sulphide,  which  blackens  a  moist  silver  coin. 

Fuming  Sulphuric  Acid,  H2S2O7,  is  made  by  adding  sulphur  trioxide 
to  sulphuric  acid,  or  by  heating  moist  ferrous  sulphate.  This  is  the 
acid  called  sulphuric  acid  by  the  alchemists.  It  is  sometimes  called 


252  Descriptive  Chemistry. 

Nordhausen  sulphuric  acid.  It  is  a  thick,  brown  liquid,  which  fumes 
strongly  in  the  air,  owing  to  the  escape  of  oxides  of  sulphur.  It  is  used 
in  gas  analysis  to  absorb  ethylene  and  other  illuminants,  and  in  dyeing 
to  dissolve  indigo.  If  the  fuming  acid  is  cooled  to  o°  C,  crystals  sepa- 
rate ;  they  are  called  pyrosulphuric  acid. 

Sodium  Thiosulphate,  Na2S2O3,  is  a  salt  of  an  unstable  acid.  It  is 
sometimes  incorrectly  called  sodium  hyposulphite,  or  simply  "  hypo." 
It  is  a  white,  crystallized  solid,  very  soluble  in  water.  The  solution, 
used  in  excess,  dissolves  the  halogen  compounds  of  silver ;  hence  its 
extensive  use  in  photography  (see  Photography).  It  also  finds  some 
use  as  an  antichlor,  and  in  chemical  analysis  for  determining  the  amount 
of  free  iodine  in  a  solution. 

Carbon  Disulphide,  CS2,  when  pure,  is  a  clear,  colorless  liquid,  with 
an  agreeable  odor.  The  commercial  substance  is  yellow  and  has  an 
offensive  odor.  It  is  poisonous.  It  is  volatile  and  extremely  inflam- 
mable, the  equation  for  its  combustion  being— - 

CS2  +    302    =  C02  +  2  SO, 

Carbon  Disulphide     Oxygen     Carbon  Dioxide      Sulphur  Dioxide 

This  liquid  is  insoluble  in  water.  It  dissolves  rubber,  gums,  fats,  resins, 
iodine,  camphor,  and  some  forms  of  sulphur.  It  is  a  highly  refracting 
liquid,  and  hollow  glass  prisms  filled  with  it  are  used  to  decompose 
light.  As  a  solvent  it  is  used  to  dissolve  pure  rubber  in  the  manufac- 
ture of  rubber  cement.  It  is  also  used  to  kill  insects  on  both  living  and 
dried  plants  (e,g.  in  museums),  and  to  exterminate  burrowing  animals, 
such  as  moles  and  woodchucks.  Many  oils,  waxes,  and  greases  are  ex- 
tracted by  carbon  disulphide.  It  is  also  used  to  manufacture  compounds 
of  sulphur  and  of  carbon. 

Until  recently  carbon  disulphide  was  manufactured  by  passing  sul- 
phur vapor  over  red-hot  coke  or  charcoal  in  iron  or  earthenware  retorts  ; 
the  product  required  laborious  purification.  It  is  now  manufactured  by 
an  electrothermal  process.  Several  groups  of  carbon  electrodes  are  set 
into  the  base  of  a  furnace,  coke  is  packed  loosely  around  them,  and  the 
body  of  the  furnace  is  filled  with  charcoal.  Sulphur  is  introduced  at 
suitable  points,  and  when  the  current  passes  the  sulphur  melts,  vapor- 
izes, and  unites  with  the  heated  carbon  above  the  electrodes. 

Selenium  and  Tellurium  are  rare  elements  which  form  compounds 
analogous  to  the  principal  compounds  of  sulphur.  These  three  with 
oxygen  form  a  natural  group,  their  physical  properties  varying  gradually 
with  increasing  atomic  weight. 


Sulphur  and  its  Compounds.  253 

EXERCISES. 

1.  What  is  the  symbol  and  atomic  weight  of  sulphur? 

2.  Where  is  free  sulphur  found?     Discuss  its  formation.     In  what 
forms  is  combined  sulphur  found?     Name  five  native  compounds  of 
sulphur.     What  animal  and  vegetable  compounds  contain  sulphur? 

3.  Give  a  brief  account  of  the  sulphur  industry  in  Sicily.     How  is 
sulphur  purified? 

4.  What  is  (a}  flowers  of  sulphur,  (<£)  brimstone,  (^)  roll  sulphur, 
(a)  milk  of  sulphur? 

5.  Summarize  the  properties  of  sulphur,  especially  its  action  when 
heated. 

6.  Describe  the  different  forms  of  sulphur. 

7.  For  what  is  sulphur  used? 

8.  What  is  hydrogen  sulphide?     Where  is  it  found?     Describe  its 
preparation. 

9.  Summarize  the  properties  of  hydrogen   sulphide.      State  the 
equation  for  its  combustion.     What  is  its  action  upon  sulphuric  acid?^ 
What  is  hydrogen  sulphide  water? 

10.    Why  is  H2S  the  formula  of  hydrogen  sulpMde2__ ^ 
n.   What  are  sulphides?     How  are  they  formed?     Name  and  de- 
scribe five.     Why  does  silverware  often  blacken?     What  use  is  made 
of  sulphides  in  qualitative  analysis? 

12.  What  is  sulphur  dioxide?     How  is  it  formed?     State  one  equa- 
tion for  its  formation.     Describe  its  preparation.     For  what  is  it  used? 

13.  Summarize  the  properties  of  sulphur  dioxide. 

14.  Why  is  SO2  the  formula  of  sulphur  dioxide? 

15.  What  is  the  volumetric  equation  for  the  formation  of  sulphur 
dioxide  from  sulphur  and  oxygen?     How  many  liters  of  oxygen  are 
needed  to  form  5  1.  of  sulphur  dioxide? 

1 6.  Discuss  sulphurous  acid  and  sulphites.     , 

17.  What  is  sulphur  trioxide?     How  is  it  prepared?     State  its  chief 
properties.     What  is  its  formula?     Why? 

1 8.  Give  a  brief  historical  account  of  sulphuric  acid.    Why  is  it  often 
called  oil  of  vitriol  ?     What  is  (a)  chamber  acid,  (b)  Nordhausen  acid, 
(c}  fuming  sulphuric  acid,  (d)  pyrosulphuric  acid? 

19.  Upon  what  fact  is  the  manufacture  of  sulphuric  acid  based?     In 
what  two  general  ways  is  the  operation  accomplished  ? 

20.  Describe  the  older  method  of  manufacturing  sulphuric  acid. 


254  Descriptive  Chemistry. 

21.  Describe  the  contact  method  of  manufacturing  sulphuric  acid. 

22.  Define  (a)  catalysis  and  (b}  catalyzer. 

23.  Summarize  the  properties  of  sulphuric  acid. 

24.  Enumerate  the  important  uses  of  sulphuric  acid. 

25.  Define  and  illustrate  (a)  sulphate,  (£)  normal  sulphate,  (<:)  acid 
sulphate. 

26.  What  is  (a)  gypsum,  (b)  white  vitriol,  (c)  green  vitriol,  (d)  blue 
vitriol,  (/)  Glauber's  salt,  (_/")  kieserite? 

27.  Describe  the  test  for  (a)  sulphuric  acid,  (b}  sulphurous  acid, 
(c)  a  soluble  sulphate,  (d)  an  insoluble  sulphate,  (tf)  a  sulphite. 

28.  State  (a)  the  properties,  and  ($)  the  uses  of  sodium  thiosulphate. 
What  is  its  common  name? 

29.  State  (a}  the  properties,  and  (b)  the  uses  of  carbon  disulphide. 
How  is  it  manufactured? 

PROBLEMS 

1.  Calculate  the  percentage  composition  of  (a)  barium   sulphate 
(BaSO4),  (£)  zinc  sulphate  (ZnSO4),  (c)  sodium  sulphate  (Na,SO4). 

2.  Calculate  the  percentage  composition  of  (a)  galena  (PbS),  (b) 
zinc  blende  (ZnS),  (c)  iron  pyrites  (FeS2),  (d)  ferrous  sulphide  (FeS). 

3.  What  weight  and  what  volume  of  hydrogen  can  be  obtained  from 
102  gm.  of  hydrogen  sulphide  ? 

4.  What  is  the  weight  of  a  stick  of  brimstone  10  cm.  long  and  4 
cm.  in  diameter  ? 

5.  How  many  grams  of  ferrous  sulphide  are  needed  to  prepare  a  liter 
of  hydrogen  sulphide  gas  ? 

6.  Sulphuric  acid  is  i  .8  times  heavier  than  water.    How  many  grams 
of  acid  will  a  liter  flask  hold  ? 

7.  Calculate  the  weight  of  oxygen  necessary  to  burn  (to  sulphur  di- 
oxide) 731  gm.  of  sulphur  containing  15  per  cent  of  impurities. 

8.  A  lump  of  sulphur  weighing  32  gm.  is  burned  in  air.     Calculate 
(a)  the  weight  of  oxygen  required,  and  (£)  the  weight  of  sulphur  di- 
oxide formed. 

9.  How  many  liters  of  oxygen  are  needed  (a)  to  form  10  1.  of 
sulphur  dioxide  by  burning  sulphur  in  air,  and  (£)  to  change  10  1.  of 
sulphur  dioxide  to  sulphur  trioxide  ? 


CHAPTER   XVIII. 
SILICON  AND  BORON. 

Occurrence  of  Silicon.  —  Silicon  does  not  occur  free  in 
nature,  being  found  almost  exclusively  as  silicon  dioxide 
(SiO2)  or  as  silicates.  These  compounds  are  so  abundant 
and  widely  distributed  that  approximately  one  fourth  of  the 
earth's  crust  is  silicon.  Sand  and  the  different  varieties  of 
quartz  are  silicon  dioxide.  Most  rocks  are  silicates. 

Silicon  itself  is  a  rare  element.  It  is  obtained  with  difficulty  by 
heating  silicon  dioxide  with  carbon,  aluminium,  or  magnesium  in  the 
electric  furnace,  or  by  heating  silicon  chloride  with  sodium. 

Like  carbon,  silicon  has  three  allotropic  forms,  —  a  brown  amorphous 
powder,  a  dark  grayish  mass  like  graphite,  and  steel-colored  crystals. 
Amorphous  silicon  may  be  changed  into  the  other  forms.  They  have 
different  properties. 

The  name  "  silicon  "  comes  from  the  Latin  word  silex,  silicis,  flint. 

Silicon  Dioxide  or  Silica,  SiO2,  is  the  most  common  com- 
pound of  silicon.  Sand,  gravel,  sandstone,  and  quartzite 
are  almost  wholly  silica.  It  is  an  essential  ingredient  of 
many  rocks,  as  granite  and  gneiss.  Quartz  is  silicon  di- 
oxide. It  has  many  varieties,  which  differ  in  color  and 
structure,  due  to  minute  impurities  or  to  the  mode  of 
formation.  Among  the  crystalline  varieties  are  the  clear, 
colorless  rock  crystal,  the  purple  amethyst,  and  the  rose, 
yellow,  glassy,  milky,  and  smoky  forms.  Varieties  imper- 
fectly crystalline  or  amorphous  are  the  waxlike  chalcedony, 
the  various  forms  of  agate  having  different  colored  layers, 
the  reddish  brown  carnelian,  the  black  and  white  onyx,  the 

255 


Descriptive  Chemistry. 


red  or  brown  jasper,  the  dull  brown  or  black  flint,  and  the 
brittle  chert.  Opal  is  hydrated  silica  (SiO2  •  nH2O).  Petri- 
fied or  silicified  wood  is  largely  some  variety  of  quartz 
which  has  replaced  the  woody  fiber.  There  is  a  "  petrified 
forest "  in  Arizona.  Infusorial  or  diatomaceous  earth  is  a 
variety  of  silica  consisting  of  the  shells  of  minute  organisms 
called  diatoms  (¥\g.  51).  Quartz  is  often  found  as  crystals 
which  consist  usually  of  a  six-sided  prism  with  a  six-sided 

pyramid  at  one  or  both 
ends,  but  the  crystals 
are  sometimes  complex 
(Fig.  52). 

Quartz     crystals     and 


FiG.  51.  —  Earth  from  Richmond,  Va.,  con- 
taining diatoms. 


FIG.  52.  —  Quartz  crystals. 


varieties  like  them  are  hard  enough  to  scratch  glass. 
They  are  insoluble  in  water  and  acids,  except  hydro- 
fluoric acid,  but  are  soluble  in  melted  hydroxides  and 
carbonates  of  sodium  and  potassium.  Quartz  is  infusible, 
except  in  the  oxyhydrogen  flame.  If  fused  with  certain 
precautions,  the  molten  mass  can  be  drawn  out  into  elastic 
threads,  which  are  used  to  suspend  delicate  parts  of  elec- 
trical instruments. 

Sandstone  and  quartzite  are  used  as  building  stones,  and 
hard  sandstone  is  made  into  grindstones  and  whetstones. 
Sand  is  used  in  making  sandpaper,  glass,  porcelain,  and 


Silicon  and  Boron.  257 

mortar.  Glass  is  roughened  and  cut  by  blowing  or  "  blast- 
ing "  fine  sand  against  it.  Many  of  the  varieties  of  quartz 
are  cut  and  polished  into  ornaments  and  gems,  e.g.  amethyst, 
opal,  and  agate.  Rock  crystal  is  used  as  the  "  diamond  " 
in  cheap  jewelry,  and  is  cut  into  lenses  for  eyeglasses  and 
optical  instruments.  Petrified  wood  is  cut  and  polished  into 
table  tops,  mantelpieces,  and  fireplaces.  Infusorial  earth  is 
used  to  polish  silver,  "electro-silicon"  being  the  commercial 
name  of  one  kind,  and  in  making  cement,  "  soluble  glass," 
dynamite,  and  refractory  brick.  Over  1300  tons  are 
annually  used  in  the  United  States. 

Silica  and  Plants.  —  Ashes  of  many  plants  contain  silica,  showing 
that  some  compound  of  silicon  is  assimilated  by  the  plant  from  the  soil 
—  probably  silicic  acid  or  a  soluble  silicate  (see  below).  The  ashes  of 
rye  and  wheat  straws  and  of  potato  stems  contain  from  40  to  70 
per  cent  of  silica.  Plants  like  horsetail,  sword  grass,  and  bamboo  are 
rich  in  silica.  The  silica  is  probably  not  a  plant  food  in  the  strict  sense, 
but  gives  firmness  to  the  tall  stalks,  especially  to  their  joints,  and  pro- 
duces the  tough  exterior  coating,  as  on  the  bamboo.  The  quills  of 
feathers  and  the  spikes  of  sponges  are  tough  and  rigid  from  the  silica 
they  contain. 

Silicon  Tetrafluoride  (SiF4)  is  formed  by  the  interaction  of  silicon 
dioxide  and  hydrofluoric  acid,  as  described  under  etching  (see  Etching). 

Silicic  Acid  and  Silicates.  —  When  silicon  dioxide  is 
fused  with  sodium  or  potassium  carbonates,  the  correspond- 
ing silicate  is  formed  thus  — 

SiO2       +       K2CO3       =       K2SiO3     +       CO2 

Silicon  Potassium  Potassium  Carbon 

Dioxide  Carbonate  Silicate  Dioxide 

Potassium  and  sodium  silicates  dissolve  in  water,  and  when 
hydrochloric  acid  is  added,  the  gelatinous  precipitate 
formed  is  a  silicic  acid  having  the  formula  H2SiO3  (proba- 


258  Descriptive  Chemistry. 

bly).  This  acid  is  decomposed,  by  heating,  into  silicon 
dioxide  and  water,  thus  — 

H2SiO3       =  SiO2         +     H2O 

Silicic  Acid  Silicon  Dioxide          Water 

There  are  many  complex  silicic  acids.  Silicates  are  salts 
of  silicic  acids,  though  they  are  often  so  complex  that  no 
actual  corresponding  acid  is  known.  Silicates  make  up 
a  large  part  of  the  earth's  crust,  silicates  of  aluminium, 
iron,  calcium,  potassium,  sodium,  and  magnesium  being  the 
most  abundant.  Many  common  rocks  and  minerals  are 
silicates,  e.g.  feldspar,  mica,  mica  schist,  hornblende,  clay, 
slate,  beryl,  garnet,  serpentine,  and  talc. 

Sodium  and  potassium  silicates  are  the  only  ones  soluble 
in  water,  and  the  thick,  sirupy  solution  is  often  called 
"water  glass  "  or  soluble  silica.  It  is  used  in  making  yel- 
low soaps,  cements,  and  artificial  stone,  to  fix  colors  in 
frescoing  and  calico  printing,  and  to  render  cloth,  wood, 
and  paper  fireproof. 

Some  forms  of  silica  dissolve  in  a  hot  solution  of  sodium  carbonate. 
Hence,  many  hot  springs,  as  in  the  Yellowstone  Park,  contain  silica  in 
solution  (as  an  alkaline  silicate),  and  when  the  water  comes  to  the  sur- 
face and  cools,  silica  is  deposited  around  the  spring  in  beautiful  forms 
called  geyserite  or  siliceous  sinter.  Probably  the  formation  of  petri- 
fied wood  is  due  to  the  deposition  of  silica  from  such  a  solution. 

Silicides  are  compounds  of  silicon  and  other  elements.  Carborun- 
dum, carbon  silicide  (or  silicon  carbide,  CSi),  has  been  mentioned  (see 
Carborundum,  Chapter  X).  Silicides  of  iron,  chromium,  and  copper 
(Fe2Si,  Cr2Si,  and  Cu2Si)  are  also  commercially  important. 

Glass  is  a  mixture  of  silicates,  one  of  which  is  always 
a  silicate  of  potassium  or  sodium.  Window  glass  is  a  sili- 
cate of  sodium  and  calcium,  and  Bohemian  glass  is  a 
silicate  of  potassium  and  calcium.  In  flint  glass,  calcium 
is  replaced  by  lead. 


Silicon  and   Boron. 


259 


Glass  is  not  made  by  mixing  silicates,  but  by  melting 
together  sand,  an  alkali,  and  a  calcium  or  a  lead  compound. 
The  alkali  may  be  sodium  carbonate  (Na2CO3),  or  potas- 
sium carbonate  (K2CO3),  or  a  mixture  of  these;  sodium 
sulphate  is  often  used.  The  calcium  compound  used  is  cal- 
cium carbonate  (CaCO3)  in  the  form  of  chalk  or  limestone. 
The  lead  compound  used  is  litharge  (PbO)  or  red  lead 
(Pb3O4).  Small  quantities  of  other  substances  are  also 
used,  e.g.  broken  glass  to  help  lower  the  melting  point  of 
the  mixture,  oxide  of  arsenic  (As2O3),  potassium  nitrate 
(KNO3),  or  manganese  dioxide  (MnO2)  to  remove  the 
greenish  color  caused  by  iron  compounds,  metallic  oxides 
or  other  substances  to -produce  colored  glass,  and  numerous 
ingredients,  such  as  calcium  fluoride  or  calcium  phosphate, 
to  make  special  kinds  of  glass. 

The  process  consists  in  heating  the  proper  mixture  in  a 
fire-clay  pot  to  a  high  temperature.  During  the  melting, 
gases  escape,  and  the  impurities,  which  rise  to  the  surface 
as  a  scum,  are  removed.  The  molten  mass  is  alldwed  to 
cool  until  it  becomes  pasty.  In  this  condition  it  may  be 
blown,  welded,  cut,  drawn,  or  molded  into  almost  any 
desired  shape. 

The  mixture  used  varies  with  the  kind  of  glass  to  be  made.  A  typi- 
cal mixture  for  table  and  bottle  glass,  used  in  a  large  works,  is  — 

Sand             1550  Ib. 

Sodium  carbonate         .         .        .         .         .'  550  Ib. 

Lime             200  Ib. 

Sodium  nitrate      .......  100  Ib. 

Total  charge 2400  Ib. 

Window  Glass  is  made  by  blowing  a  lump  of  glass  into  a  hollow 
globe  and  then  into  a  cylinder ;  this  on  being  opened  at  both  ends  and 
cut  lengthwise  spreads  open  flat.  Plate  glass,  which  has  about  the 


260  Descriptive  Chemistry. 

same  composition  as  window  glass,  is  made  by  pouring  the  molten 
glass  upon  a  large  table,  rolling  it  with  a  hot  iron  roller,  and  subse- 
quently grinding  and  polishing  it.  Plate  glass  is  used  for  large  win- 
dows and  for  mirrors,  but  considerable  rough  plate  is  used  for  skylights 
and  floors.  Crown  glass  is  a  good  quality  of  window  glass.  It  has  a 
brilliant  surface.  Limited  quantities  are  used  as  "  bull's  eyes  "  in  deco- 
rative windows.  Bohemian  glass  is  the  hard  glass  of  which  much 
chemical  apparatus  is  made.  Flint  glass  is  a  silicate  of  potassium  and 
lead ;  it  is  a  lustrous,  soft  glass,  largely  used  in  making  lamp  chimneys 
and  globes.  Pure  flint  glass  is  often  called  strass  or  paste,  and  on  ac- 
count of  its  luster  and  brilliancy  it  is  made  into  artificial  gems.  Lenses 
for  telescopes  and  other  optical  instruments  usually  consist  of  both 
crown  and  flint  glass.  Cut  glass  is  flint  glass.  The  object  is  first 
molded  or  blown  into  the  general  shape,  the  design  is  then  cut  into 
the  soft  glass  by  a  wheel,  and  the  finished  object  is  polished  by  a 
wooden  wheel  smeared  with  rouge  (oxide  of  iron)  or  putty. 

Many  objects,  such  as  tumblers  and  small  dishes,  are  now  made  by 
pressing  the  soft  glass  with  a  die  or  by  blowing  it  into  a  mold.  Fruit 
jars,  bottles,  and  lamp  chimneys  are  blown  by  machinery.  Many  other 
improvements  have  increased  the  output  and  improved  the  quality  of 
glass. 

All  glassware  must  be  cooled  slowly  to  prevent  the  glass  from  being 
brittle.  This  operation  is  called  annealing,  and  is  accomplished  by 
passing  the  objects  slowly  through  a  furnace  in  which  the  temperature 
is  gradually  lowered. 

Glass  is  colored  by  adding  different  substances  which  dissolve  in  the 
molten  mass.  Iron  and  chromium  compounds  make  it  green,  the  green 
color  of  many  bottles  and  fruit  jars  being  due  to  the  iron  in  the  cheap 
materials  used ;  copper  and  cobalt  compounds  produce  different  shades 
of  blue ;  manganese  dioxide  gives  a  pink  or  a  violet,  and  a  mixture  of 
manganese  dioxide  and  iron  oxide  gives  an  orange  color ;  yellow  is  pro- 
duced by  charcoal,  sulphur,  or  silver;  certain  copper  compounds  or 
gold  give  a  ruby  color ;  translucent  or  white  glass  is  made  by  adding 
fluor  spar  or  cryolite ;  smoked  glass  contains  nickel ;  iridescent  glass 
is  made  by  exposing  it  to  the  vapors  of  hydrochloric  acid  or  of  tin 
chloride  (SnCl4). 

The  United  States  produces  yearly  about  50,000,000  dollars'  worth 
of  glass.  The  industry  is  carried  on  in  about  twenty-five  states,  Penn- 
sylvania producing  two  fifths  of  the  total  output. 


Silicon  and  Boron.  261 


BORON. 

Occurrence. —  Boron  is  never  found  free,  but  the  com- 
pounds, borax  (Na2B4O7)  and  boric  acid  (H3BO3),  are 
abundant. 

Boron  itself  is  an  uncommon  element.  It  is  prepared  by  heating 
the  oxide  (B2O3)  with  magnesium,  aluminium,  sodium,  or  potassium. 
It  is  greenish  brown  amorphous  powder,  without  taste  or  odor.  It 
burns  when  heated  in  air,  forming  the  oxide  (B2O3).  It  also  unites 
with  the  halogens,  sulphur,  and  nitrogen.  It  forms  many  borides,  one 
of  which,  carbon  boride  (CB,.),  is  said  to  be  harder  than  diamond. 

Boric  Acid,  H3BO3,  is  contained  in  the  waters  and  steam 
of  certain  volcanic  regions,  notably  Tuscany.  Large 
basins  or  tanks  are  built  around  these  steam  jets,  and 
are  arranged  so  that  the  water  flows  at  intervals  from  one 
reservoir  into  the  next  lower,  constantly  becoming  charged 
with  more  boric  acid,  as  the  steam  condenses.  The  final 
solution  is  evaporated  by  aid  of  the  heat  from  the  steam 
jets,  and  the  crude  boric  acid  ;which  settles  out  is  purified 
by  recrystallization.  This  compound  is  sometimes  called 
boracic  acid. 

Considerable  boric  acid  is  also  made  in  California  from  borax,  and 
in  Germany  from  the  boracite  found  at  Stassfurt. 

Boric  acid  crystallizes  in  lustrous,  white  flakes,  which  feel  greasy. 
It  dissolves  slightly  in  cold  water,  readily  in  hot  water,  and  in  alcohol. 
When  the  alcoholic  solution  is  burned,  a  boron  compound  colors  the 
vapor  green.  This  is  the  test  for  boron  compounds. 

Boric  acid  is  used  in  making  borax,  in  the  manufacture  of  enamels 
and  glazes  for  pottery,  as  an  antiseptic  in  medicine  and  surgery,  and  for 
preserving  meat,  fish,  milk,  butter,  beer,  and  wine. 

Borax,  Na2B4O7.  ioH2O,  occurs  in  large  quantities  in 
California,  and  an  impure  borax  called  tinkal  comes  from 
Tibet.  Much  of  the  commercial  borax  is  made  from 


262 


Descriptive  Chemistry. 


boric  acid  or  from  native  calcium  borate  (colemanite, 
Ca2B6On .  5  H2O)  by  boiling  with  sodium  carbonate  and 
separating  the  borax  by  crystallization. 

Borax  is  a  white  crystallized  solid,  having  ten  or  five 
molecules  of  water  of  crystallization.  It  effloresces  in  the 
air.  When  heated,  ordinary  borax  melts,  then  swells  up 
into  a  white  porous  mass,  which  finally  becomes  a  glassy 
solid.  This  glassy  borax  dissolves  metallic  substances, 
especially  oxides.  If  the  borax  is  melted  on  the  end  of 
a  looped  platinum  wire,  the  transparent  globule  is  called  a 
borax  bead.  These  beads  differ  in  color  under  .different 
circumstances,  and  the  oxides  of  metals  cause  the  beads  to 
assume  colors  which  are  characteristic  of  the  metals,  as 
may  be  seen  by  the  following  table  :  — 


COLORS  OF  BORAX  BEADS. 


OXIDIZING  FLAME. 

REDUCING  FLAME. 

METAL 

Hot. 

Cold. 

Hot. 

Cold. 

Chromium     . 

Reddish  yellow 

Yellowish  green 

Green 

Green 

Cobalt      .     . 

Blue 

Blue 

Blue 

Blue 

Copper     .     . 

Green 

Greenish  Blue 

Colorless 

Red 

Manganese  . 

Violet 

Violet 

Colorless 

Colorless 

The  bead  test  is  often  used  in  chemistry  to  confirm  other 
observations  or  to  suggest  further  examination. 

Borax  is  used  in  the  manufacture  of  enamels  and  glazes, 
and  in  the  formation  of  the  "  paste  "  for  artificial  gems. 
Immense  quantities  are  used  for  preserving  canned  meat 
and  fish.  It  is  a  cleansing  agent,  and  large  quantities  are 
consumed  in  laundries  as  well  as  in  the  manufacture  of 


Silicon  and  Boron.  263 

soaps,  particularly  those  intended  for  use  in  hard  water 
(see  Soap).  Its  power  to  dissolve  oxides  adapts  it  for  use 
in  soldering  metals.  Solder  adheres  only  to  clean  metals, 
so  a  little  borax  is  used  to  dissolve  the  film  of  oxide  on 
the  surfaces  to  be  joined.  It  is  likewise  used  in  welding 
metals  and  as  a  flux  in  their  preparation.  Considerable 
quantities  are  used  as  a  mordant  in  calico  printing  and  in 
dyeing.  It  is  an  ingredient  of  ointments,  lotions,  and 
powders,  which  are  designed  to  relieve  hoarseness  or 
skin  eruption. 

EXERCISES. 

1.  What  is  the  symbol   and   atomic   weight  of  (a)   silicon,  and 
(fr)  boron  ? 

2.  How  is  silicon  found  in  nature  ?     What  proportion  of  the  earth's 
crust  is  combined  silicon  ? 

3.  Name  several  common  forms  of  silicon  dioxide.     Describe  the 
different  varieties  of  quartz. 

4.  What  is  (a)  petrified  wood,  (£)  opal,  (c)  diatomaceous  earth, 
(d)  "  electro-silicon  "  ? 

5.  Summarize  the  properties  of  quartz.     How  can  it  be  readily 
distinguished  from  other  minerals  and  rocks  ? 

6.  State  the  uses  of  the  different  forms  of  silicon  dioxide. 

7.  Discuss  the  relation  of  silicon  dioxide  to  plants. 

8.  Review  with  special  reference  to  silicon  compounds   (a)  car- 
borundum, and  (£)  etching  glass. 

9.  Describe  the  formation  and  state  the  properties   of  ordinary 
silicic  acid.     Name  several  common  silicates.     What  metals  are  com- 
ponents of  silicates  ? 

10.  Describe  the  formation,  state  the  uses,  and  enumerate  the  prop- 
erties of  "water  glass." 

11.  What  is  glass  ?    How  is  it  made  ?    Name  the  components  of 
the  different  kinds. 

12.  What  is  (a)  window  glass,  (£)  plate  glass,  (c)  Bohemian  glass, 
(d)  flint  glass,  and  (e)  cut  glass  ? 

13.  How  is  glass  (#)  annealed,  and  ($)  colored  ? 

14.  How  is  boron  found  in  nature  ?    What  is  the  formula  of  (#) 
borax,  and  (£)  boric  acid  ? 


264  '    Descriptive  Chemistry. 

15.  Where  is  boric  acid  found  ?    How  is  it  manufactured  ?    State 
its  properties  and  uses. 

1 6.  Where  is  borax  found  ?    How  is  it  prepared  for  commerce  ? 
State  its  properties  and  uses. 

17.  Describe  the  borax  bead.     State  and  illustrate  its  use. 


PROBLEMS. 

1 .  Calculate  the  percentage  composition,  of  (#)  willemite  (Zn2SiO4), 
(b)  steatite  (MgsSi4OH.  H2O),  (V)  quartz  (SiO2). 

2.  What  per  cent  of  borax  (Na2B4O7.  10  H2O)  is  boron  ? 


CHAPTER  XIX. 
PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH. 

PHOSPHORUS,  arsenic,  antimony,  and  bismuth,  together 
with  nitrogen,  form  a  natural  group  of  elements. 

PHOSPHORUS. 

Occurrence.  —  Free  phosphorus  is  not  found  in  nature, 
but  phosphates  are  numerous  and  abundant.  The  most 
common  are  phosphorite  (impure  Ca3(PO4)2)  and  apatite 
(3  Ca3(PO4)2.CaCl2  or  CaF2).  About  o.i  per  cent  of  the 
earth's  crust  is  phosphorus.  Calcium  phosphate  is  pres- 
ent in  all  fertile  soils,  being  a  product  of  decayed  rocks. 
Plants  and  animals  contain  phosphorus  compounds  as 
essential  constituents  of  the  brain,  nerves,  and  bones. 

Phosphorus  was  discovered  in  1669  by  Brand,  who  obtained  it  by 
heating  a  certain 'kind  of  animal  matter.  Scheele,  in  1771,  extracted  it 
from  bones. 

Preparation.  —  Phosphorus  is  too  dangerous  a  substance  to  prepare 
in  the  laboratory,  (i)  It  is  manufactured  from  bone  ash  or  from  native 
phosphates.  The  finely  ground  material  is  mixed  in  large  vats  with 
enough  sulphuric  acid  to  produce  the  following  change:  — 

Ca,(P04)2     +         3H,S04         =         2H3PO4  +     3  CaSO4 

Calcium  Sulphuric  Acid         Phosphoric  Acid  Calcium 

Phosphate  (Ortho-)  Sulphate 

The  insoluble  calcium  sulphate  is  removed  by  filtering  the  mixture 
through  cinders.  The  phosphoric  acid  solution  is  concentrated,  mixed 
with  sawdust,  coke,  or  charcoal,  and  dried,  being  changed  thereby 
according  to  the  equation  — 

H,PO4  .  HPO3  +         H2O 

Phosphoric  Acid  (Ortho-)         Phosphoric  Acid  (Meta-) 

265 


266 


Descriptive  Chemistry. 


The  dried  mass  is  heated  to  a  high  temperature  in  clay  retorts  arranged 
in  tiers  (Fig.  53),  the  change  thus  produced  being  substantially  — 


2H2         +  12  CO 
Hydrogen        Carbon 
Monoxide 


4HPO3        +     12  C       =  P4         + 

Phosphoric  Acid    Carbon         Phosphorus 
(Mela-) 

The  phosphorus  distils  as  a  vapor  through  a  pipe  into  a  trough  of  water, 
where  it  collects  as  a  heavy  liquid.  (2)  Phosphorus  is  also  manufactured 
in  the  electric  furnace.  A  mixture  of  a  phosphate,  carbon,  and  sand  is 
fed  into  a  furnace  provided  with  an  outlet  pipe  through  which  the  phos- 
phorus vapor  passes  into  a  condenser.  The  residue  is  drawn  off  as  a 
slag  at  the  bottom.  The  equation  for  the  chemical  change  is  — 


2Ca3(P04)2 

Calcium 
Phosphate 


6SiO2     +        10  C   =        P4         +    10  CO  +  3CaSiO3 
Sand  Carbon    Phosphorus     Carbon        Calcium 

Monoxide       Silicate 


Either  method  gives  a  black  product,  which  is  purified  by  redistil- 
lation in  an  iron  retort,  or  by  oxidation  under  water  with   sulphuric 

acid  and  potassium  dichromate ; 
finally  it  is  pressed  through  can- 
vas bags  and  molded  into  sticks. 

Properties.  — Phosphor- 
us has  three  allotropic 
modifications,  —  yellow  or 
ordinary,  red  or  amorphous, 
and  black  or  crystalline. 
Ordinary  phosphorus  is 
a  yellowish,  translucent 
solid.  The  color  deepens 
by  exposure  to  light.  At 
ordinary  temperatures 
phosphorus  is  like  wax, 
but  at  low  temperatures  it 
is  brittle.  Under  water  it 
melts  at  44°  C.  Exposed 

FIG.  53.  —  Apparatus  for  the  manufacture 

of  phosphorus.  to  the  air   it   immediately 


Phosphorus,  Arsenic,  Antimony,  Bismuth.     267 

gives  off  white  fumes,  and  at  34°  C.  takes  fire  and  burns 
with  a  brilliant  flame,  the  main  product  being  phosphorus 
pentoxide  (P2O5).  In  moist  air  it  glows,  as  may  be  easily 
seen  by  rubbing  the  head  of  a  match  in  a  dark  room. 
This  property  gave  the  element  its  name  (from  the  Greek 
word  phosphoros,  light  bringer).  The  ease  with  which  it 
ignites  makes  phosphorus  dangerous  to  handle.  Burns 
from  it  are  severe  and  hard  to  heal.  It  is  very  poisonous, 
and  the  workmen  in  phosphorus  factories  are  subject  to 
a  dreadful  disease,  which  rots  the  bones.  A  fatal  dose 
is  about  o.  1 5  gm.  Phosphorus  is  kept  beneath  water,  and 
should  never  be  handled  or  cut  unless  so  covered.  It  is 
nearly  insoluble  in  water,  but  dissolves  in  carbon  disulphide 
and  slightly  in  sodium  hydroxide  solution.  .Yellow  phos- 
phorus has  a  faint  odor,  which  may  be  easily  detected  by 
smelling  a  match  head.  Red  phosphorus  is  made  by 
heating  ordinary  phosphorus  to  25O°-3OO°  C.  in  a  closed 
vessel.  Any  unchanged  yellow  phosphorus  is  extracted 
with  sodium  hydroxide  solution.  The  red  phosphorus  is 
usually  a  reddish  brown  powder,  though  sometimes  it  is 
a  brittle  mass.  It  is  opaque  and  odorless,  does  not  give 
light,  nor  can  it  be  easily  ignited.  It  is  poisonous,  and 
does  not  dissolve  in  carbon  disulphide.  Its  specific  gravity 
is  2.25,  that  of  the  yellow  form  being  1.836.  It  can  be 
handled  without  danger.  Heated  to  about  260°  C.  in  an 
atmosphere  of  nitrogen  or  carbon  dioxide,  it  changes  into 
ordinary  phosphorus. 

Black  Phosphorus  is  formed  by  dissolving  red  phosphorus  in  melted 
lead,  and  allowing  crystals  to  separate.  Its  specific  gravity  is  2.34. 

The  vapor  density  of  phosphorus  is  such  that  its  molecule  must 
contain  four  atoms,  hence  its  molecular  formula  is  P4. 

Certain  rat  and  bug  poisons  contain  ordinary  phosphorus,  but  most 
of  the  phosphorus  of  commerce  is  consumed  in  the  manufacture  of 
matches  (see  below). 


268  Descriptive  Chemistry. 

Oxides  of  Phosphorus.  —  The  two  important  oxides  are  phosphorus 
or  trioxide  (P2O3  or  P4O6)  and  phosphoric  or  pentoxide  (P2O5).  Phos- 
phorous oxide  is  a  white  solid  formed  by  the  slow  oxidation  of  phos- 
phorus or  by  burning  phosphorus  in  a  limited  supply  of  air.  It  has  the 
odor  of  phosphorus  and  is  poisonous.  Warmed  in  the  air,  it  changes 
into  the  pentoxide.  It  unites  with  water  to  form  phosphorous  acid, 
thus-  p^  +  3H2o  =  2H,P03 

Phosphorous  Oxide  Phosphorous  Acid 

Phosphoric  Oxide  (P2O5)  is  the  white,  snowlike  solid  formed  by 
burning  phosphorus  in  an  abundant  supply  of  air.  It  is  very  deli- 
quescent, quickly  withdrawing  moisture  from  the  air  and  combining  vig- 
orously with  water  with  a  hissing  noise.  It  resembles  sulphur  trioxide 
in  its  power  to  char  wood  and  paper  by  withdrawing  from  them  the 
elements  of  water.  It  is  often  used  in  the  laboratory  to  dry  gases. 

Acids  and  Salts  of  Phosphorus.  —  There  are  three 
phosphoric  acids,  —  orthophosphoric  (H3PO4),  metaphos- 
phoric  (HPO3),  and  pyrophosphoric  (H4P2O7).  Phos- 
phorous acid  (H3PO3)  and  hypophosphorous  acid  (H3PO2) 
are  important  compounds. 

Orthophosphoric  Acid  is  a  by-product  in  the  manufacture  of  phos- 
phorus from  bone  ash  (see  above)  ;  it  may  be  made  by  oxidizing  red 
phosphorus  with  nitric  acid,  or  by  dissolving  phosphorus  pentoxide 
in  hot  water,  thus  — 

PA  +         3H20  2H3P04 

Phosphorus  Pentoxide  Orthophosphoric  Acid 

It  is  a  white,  crystalline  deliquescent  solid. 

Metaphosphoric  Acid  is  formed  by  heating  orthophosphoric  acid  to  a 
high  temperature,  thus  — 

H3P04  HP03  +  H20       / 

Orthophosphoric  Acid        Metaphosphoric  Acid 
It  may  be  formed  by  dissolving  the  pentoxide  in  cold  water,  thus  — 
P2O5  -f   H2O   =   2HPO3. 

At  ordinary  temperature  it  is  a  glassy  solid,  and  is  called  glacial  phos- 
phoric acid.  It  dissolves  readily  in  water,  and  the  solution  changes 
into  orthophosphoric  acid  —  slowly  in  the  cold,  rapidly  when  boiled. 


Phosphorus,  Arsenic,  Antimony,  Bismuth.     269 

Pyrophosphoric  Acid  is  formed  by  heating  orthophosphoric  acid  to 
200° -300°  C.,  thus  — 

2H3PO4  =  H4P2O7  +     H2O 

Orthophosphoric  Acid        Pyrophosphoric  Acid 

A  sodium  salt  of  the  ortho-acid  is  usually  used.     It  may  also  be  formed 
thuS~  PA  +   2H20   =  H4P207.  ^ 

This  acid  is  an  amorphous,  glassy  (but  sometimes  crystalline)  solid. 
It  is  readily  soluble  in  water,  and  its  solution  behaves  like  metaphos- 
phoric  acid. 

Orthophosphoric  acid  is  tribasic,  and  its  salts,  the  phosphates,  are 
numerous.  The  most  important  is  the  normal  calcium  salt,  Ca3(PO4)2. 
Hydrogen  disodium  phosphate  (HNa2PO4)  is  the  commercial  sodium 
phosphate.  This  salt  and  hydrogen  sodium  ammonium  phosphate,  or 
microcosmic  salt  (HNa(NH4)PO4),  are  used  in  chemical  analysis.  The 
"  acid  phosphate  "  sold  as  a  beverage  is  a  solution  of  one  or  more  acid 
calcium  phosphates  (HCaPO4  and  H4Ca(PO4)2).  Metaphosphates  are 
formed  by  heating  primary  or  (mono-)  sodium  phosphates,  thus  — 

H2NaPO4  NaPO3         +     H2O 

Primary  Sodium 

Sodium  Phosphate      Metaphosphate 

Pyrophosphates  are  formed  by  heating  secondary  (or  di-)  phosphates, 
thus  —    2HNa,P04  Na4P2O7  +     H2O 

Disodium  Phosphate      Sodium  Pyrophosphate 

Hypophosphites  are  produced  by  treating  phosphorus  with  alkalies. 
They  are  often  used  as  medicines. 

Other  Compounds  of  Phosphorus.  —  Phosphine  (PH3)  is  analogous 
to  ammonia  (NH3),  though  it  is  not  alkaline.  It  is  made  by  heating 
sodium  (or  potassium)  hydroxide  with  phosphorus.  It  is  poisonous, 
has  a  disagreeable  odor,  and  burns  in  the  air,  owing  to  the  presence  of 
an  inflammable  compound  of  phosphorus  and  hydrogen.  Phosphine 
itself  does  not  burn.  It  combines  with  other  substances,  forming 
phosphonium  compounds,  which  are  analogous  to  ammonium  com- 
pounds, e.g.  — 

PH3  +  HI  PH4I        >/ 

Phosphine  Hydriodic  Acid      Phosphonium  Iodide 


270  Descriptive  Chemistry. 

Phosphorus  Trichloride  (PCI.,)  is  a  disagreeable  smelling  liquid,  made 
by  the  combustion  of  dry  chlorine  and  phosphorus ;  and  phosphorus 
pentachloride  (PC1-)  is  a  greenish  solid  made  by  passing  chlorine  into 
a  vessel  containing  the  trichloride. 

Matches.  —  Phosphorus  is  chiefly  used  in  the  manufac- 
ture of  matches.  Soft  wood  is  cut  by  machinery  into  the 
desired  shape.  The  cards  or  sticks  are  fixed  in  a  frame, 
and  one  end  is  first  dipped  into  melted  sulphur  or  paraffin 
and  then  into  the  phosphorus  mixture.  The  latter  consists 
usually  of  different  proportions  of  phosphorus,  manganese 
dioxide,  glue,  and  a  little  coloring  matter.  Manganese  di- 
oxide may  be  replaced  by  other  oxidizing  agents.  These 
matches  are  the  ordinary  friction  or  sulphur  kind.  By 
rubbing  them  on  a  rough  surface  enough  heat  is  gener- 
ated to  cause  the  phosphorus  to  unite  with  the  oxygen  of 
the  oxidizing  agent,  and  the  heat  thereby  produced  sets 
fire  to  the  sulphur  or  paraffin,  and  this  in  turn  kindles  the 
wood.  Since  these  matches  are  poisonous,  and  liable  to 
take  fire,  their  manufacture  has  been  prohibited  in  some 
countries  (e.g.  Switzerland  and  the  Netherlands).  Safety 
matches,  which  replace  them,  contain  no  yellow  phos- 
phorus. The  head  of  this  kind  is  usually  a  colored  mix- 
ture of  antimony  sulphide,  potassium  chlorate,  and  glue; 
while  the  surface  upon  which  the  match  must  be  rubbed  to 
light  is  coated  with  a  mixture  of  red  phosphorus,  glue,  and 
powdered  glass.  Matches  are  made  by  machinery,  several 
million  being  produced  in  one  day. 

Relation  of  Phosphorus  to  Life.  —  Phosphorus  is  essen- 
tial to  the  growth  of  plants  and  animals.  Plants  take 
phosphates  from  the  soil  and  store  up  the  phosphorus 
compounds,  especially  in  their  fruits  and  seeds.  Animals 
eat  this  vegetable  matter,  assimilate  the  phosphorus  com- 
pounds, and  deposit  them  in  the  bones,  brain,  and  nerve 


Phosphorus,  Arsenic,  Antimony,  Bismuth.     271 

tissue.  Bones  contain  about  60  per  cent  of  calcium  phos- 
phate. Part  of  the  combined  phosphorus  consumed  by 
animals  is  rejected  by  them,  and  often  finds  its  way  back 
into  the  soil. 

The  constant  removal  of  phosphates  by  plants  would  soon  exhaust 
the  soil.  Hence  phosphorus  is  restored  to  the  soil  in  the  form  of  natu- 
ral or  artificial  fertilizers.  Natural  fertilizers  are  (i)  stable  refuse, 
which  always  contains  some  of  the  phosphates  from  the  food  originally 
fed  to  the  animals  ;  (2)  guano,  which  is  the  dried  excrement  and  carcasses 
of  the  sea  birds  that  once  lived  in  vast  numbers  in  Peru  and  Chili ;  and 
(3)  phosphate  slag,  which  is  a  phosphorus  by-product  obtained  in  manu- 
facturing steel.  These  and  bones  are  ground  and  spread  upon  the  soil. 
Artificial  fertilizers  are  made  from  phosphate  rock.  This  occurs  in  large 
beds  in  South  Carolina,  Tennessee,  and  Florida,  which  yield  about  a 
million  tons  a  year.  It  consists  of  the  hardened  remains  of  land  and 
marine  animals,  and  is  mainly  tricalcium  phosphate  (Ca3(PO4)9).  It  is 
insoluble  in  water,  and  must  be  changed  into  the  soluble  monocalcium 
salt  (H4Ca(PO4)2,  so  that  it  can  be  evenly  distributed  through  the  soil 
and  easily  taken  up  by  plants.  This  soluble  salt  is  called  "  superphos- 
phate of  lime."  When  phosphate  rock  is  treated  with  sulphuric  acid, 
the  changes  involved  may  be  written  thus  — 

Ca3(PO4)2      +    '2H2SO4      =      H4Ca(PO4)2      +      2CaSO4 
Tricalcium  "  Superphosphate  Calcium 

Phosphate  of  Lime "  Sulphate 

Ca3(P04)2      +      3H2S04      =         2H3P04         +      3  CaSO4 

Phosphoric  Acid 

Ca3(PO4)2      +        H2SO4       =      H2Ca2(PO4)2     +       CaSO4 

Dicalcium  Phosphate 

The  aim  is  to  convert  the  crude  phosphate  rock  into  "superphos- 
phate," but  the  three  reactions  usually  occur.  The  product  is  ground, 
dried,  and  packed  in  bags  for  the  market.  On  standing,  it  may  undergo 
"  reversion,"  i.e.  the  "  superphosphate  "  and  phosphoric  acid  may  form 
insoluble  phosphates,  thus  making  the  fertilizer  less  valuable.  Some- 
times "  superphosphate  "  is  mixed  with  compounds  of  nitrogen  and  of 
potash  to  produce  a  complete  fertilizer. 


272  Descriptive  Chemistry. 

ARSENIC. 

Occurrence.  —  Arsenic  is  found  free  in  nature,  but  it 
usually  occurs  combined  with  sulphur  or  a  metal,  or  with 
both.  The  common  arsenic  ores  are  realgar  (As2S2), 
orpiment  (As2S3),  arsenic  pyrites  or  mispickel  (FeSAs). 
Arsenic  trioxide  or  arsenolite(As2O3)  is  also  found.  Small 
quantities  of  arsenic  occur  in  many  ores. 

The  United  States  annually  imports  over  6,000,000  pounds  of  arsenic 
and  its  compounds,  mainly  from  England  and  Germany. 

Arsenic  is  prepared  in  the  laboratory  by  heating  a  mixture  of  arse- 
nious  oxide  and  charcoal  in  a  glass  tube.  The  change  is  represented 
thus  — 

2As3O8          -f          6C  As4     +  6  CO 

Arsenious  Oxide  Carbon  Arsenic      Carbon  Monoxide 

On  a  large  scale  it  is  extracted  from  its  ores  either  by  the  above  method 
or  by  roasting  arsenic  pyrites  (FeSAs)  in  the  absence  of  oxygen. 

Arsenic  has  marked  properties.  It  is  a  brittle,  steel-gray  solid.  A 
freshly  broken  piece  has  a  metallic  luster,  which  disappears  slowly  in  a 
moist  atmosphere.  It  tends  to  crystallize.  The  specific  gravity  is  from 
5.62  to  5.96.  Heated  in  the  air,  it  volatilizes  without  melting,  and  the 
vapor  has  an  odor  like  garlic.  At  about  180°  C.  it  burns  in  the  air  with 
a  bluish  flame,  forming  the  white  oxide  (As2O3).  Arsenic  molecules, 
like  those  of  phosphorus,  contain  four  atoms.  In  some  respects  arsenic 
resembles  both  metals  and  non-metals.  It  is  used  to  harden  the  lead 
which  is  made  into  shot. 

Arsenious  Oxide  or  Arsenic  Trioxide,  As2O3,  is  the 
most  important  compound  of  arsenic,  and  is  often  called 
simply  " arsenic"  or  "white  arsenic."  It  is  found  free  in 
nature,  but  is  usually  manufactured  by  roasting  arsenic 
ores.  There  are  two  common  varieties,  a  white,  granular 
powder  and  an  amorphous,  glasslike  solid.  It  has  no  odor, 
a  faint,  metallic  taste,  dissolves  slightly  in  cold  water,  but 
readily  in  hot  hydrochloric  acid.  Arsenic  trioxide  is  a 


Phosphorus,  Arsenic,  Antimony,  Bismuth.     273 

rank  poison.  The  antidote  is  fresh  ferric  hydroxide,  which 
is  made  by  adding  ammonium  hydroxide  to  a  ferric  salt, 
e.g.  ferric  chloride.  Small  doses  (2  to  3  grains)  are  usually 
fatal,  but  by  habitual  use  the  system  appropriates  larger 
doses  without  ill  effects.  Workmen  in  arsenic  factories 
often  accidentally  swallow  with  impunity  quantities  which 
would  ordinarily  prove  fatal.  It  is  used  for  making  pig- 
ments for  green  paints,  for  fly  and  rat  poison,  in  mak- 
ing glass,  arsenic  compounds,  in  calico  printing,  and  in 
preserving  skins.  As  a  medicine  it  is  used  to  purify  the 
blood. 

Other  Arsenic  Compounds. — The  native  mineral  orpiment  (As2S3) 
is  used  in  making  a-  yellow  paint,  and  realgar  (As2S2)  a  red  paint. 
Scheele's  green  is  chiefly  copper  arsenite  (HCuAsO3),  and  was  formerly 
used  to  make  a  cheap  green  paint  and  to  color  wall  paper.  The  com- 
plex arsenic  compound  Paris  green  is  a  light  green  powder ;  owing  to 
its  poisonous  character  it  is  used  to  exterminate  potato  bugs  and  other 
insects.  Arsenic  forms  acids  analogous  to  the  acids  of  phosphorus, 
though  they  are  less  important.  The  salts  sodium  arsenate  (HNa2AsO4) 
and  arsenite  (NaAsO2)  are  used  in  dyeing.  The  formation  of  the  yel- 
low sulphide  (As2S3)  by  passing  hydrogen  sulphide  into  an  arsenic 
solution  containing  hydrochloric  acid  is  the  usual  test  for  arsenic. 

Marsh's  Test  for  Arsenic.  — Arsenic  itself  is  not  poisonous,  but  its 
compounds  are  among  the  most  poisonous  substances  known.  For- 
tunately, combined  arsenic  is  easily  detected  by  a  simple  method,  called 
Marsh's  test.  An  apparatus  for  generating  hydrogen  is  provided  with 
a  hard  glass  horizontal  delivery  tube,  narrowed  in  places  and  drawn  to 
a  point.  Pure  zinc,  pure  dilute  sulphuric  acid,  and  the  arsenic  solution 
are  put  in  the  generator.  Hydrogen  and  gaseous  hydrogen  arsenide 
(or  arsine  (AsHo)  )  are  formed.  If  this  mixture  is  lighted  at  the  end 
of  the  delivery  tube,  metallic  arsenic  is  deposited  as  a  black  coating  on 
cold  porcelain  held  in  the  flame ;  or  if  the  tube  is  heated  in  front  of  a 
narrow  place,  arsenic  is  deposited  at  this  point.  This  deposit  dissolves 
in  sodium  hypochlorite  solution,  but  a  deposit  of  antimony,  similarly 
produced,  does  not  dissolve.  By  this  delicate  test  the  merest  trace  of 
arsenic  is  readily  and  positively  detected. 


274  Descriptive  Chemistry. 

ANTIMONY. 

Occurrence  of  Antimony.  —  Small  quantities  of  free  anti- 
mony are  found.  The  most  common  ore  is  stibnite  (Sb2S3), 
which  occurs  in  Japan,  Austria-Hungary,  France,  Algeria, 
Italy,  Mexico,  and  Turkey.  Large  deposits  in  California 
and  Nevada  are  now  utilized,  about  3,000,000  pounds  being 
annually  produced. 

Stibnite  was  known  in  the  fifteenth  century.  The  Latin  name  of 
antimony  is  stibium,  from  stibnite,  which  gives  the  symbol  of  the 
element,  Sb. 

Antimony  is  prepared  on  a  large  scale  by  two  methods.  In  one  the 
sulphide  is  roasted,  and  the  oxide  thus  formed  is  reduced  with  charcoal. 
Equations  representing  the  main  changes  are  — 

2Sb2S3          +         9O2        =         2SbO3         4         6  SO2 
Antimony  Sulphide         Oxygen        Antimony  Oxide     Sulphur  Dioxide 

2Sb2O3         +          3C  4Sb  +         3CO2 

The  other  method  consists  in  heating  the  sulphide  with  iron,  the  equation 
for  the  chemical  change  being  — 

Sb2S3  +         3Fe       =         2Sb          -f  3  FeS 

Antimony  Sulphide  Iron  Antimony  Iron  Sulphide 

Antimony  has  interesting  properties.  It  is  a  silver  white,  crystal- 
line, brittle  solid.  Its  specific  gravity  is  6.7.  At  ordinary  temperatures 
antimony  does  not  tarnish  in  the  air,  but  when  heated,  it  burns  with  a 
bluish  flame,  forming  the  white,  powdery  antimony  trioxide  (Sb2O.,). 
Powdered  antimony  burns  brilliantly  when  added  to  chlorine,  bromine, 
or  iodine.  Nitric  acid  oxidizes  it,  and  aqua  regia  dissolves  it.  Anti- 
mony melts  at  about  450°  C.  It  expands  on  cooling,  and  is  therefore 
one  constituent  of  type  metal  (see  Alloys  of  Lead). 

Compounds  of  Antimony.  —  Antimony  forms  stibine  (SbH3),  which 
is  analogous  to  ammonia  (NH3)  and  arsine  (AsH3),  pyro-  and  meta- 
acids,  the  oxides,  Sb2O3  and  Sb2O-,  and  halogen  compounds.  It  also 
forms  complex  compounds  in  which  antimony  acts  as  a  metal.  Tartar 
emetic  is  potassium  antimonyl  tartrate  (KSbO  .C4H4O6).  It  is  used  as 
a  medicine  and  as  a  mordant  in  dyeing  cotton.  Antimony  trisulphide 


Phosphorus,  Arsenic,  Antimony,  Bismuth.     275 

(Sb.,S3)  is  a  reddish  solid,  formed  by  passing  hydrogen  sulphide  gas 
into  a  solution  of  antimony  —  the  test  for  antimony.  The  sulphide  is 
used  in  making  the  red  rubber  tubing  and  stoppers  used  in  the  labora- 
tory. The  chlorides  (SbCl3  and  SbCl3)  are  formed  by  the  action  of 
chlorine  upon  the  metal ;  with  water  they  form  the  white  solids  called 
oxychlorides,  e.g.  SbOCl.  The  formation  of  antimony  oxychloride  is 
sometimes  used  as  a  test  for  antimony,  but  the  more  common  test  is 
the  formation  of  the  reddish  orange  sulphide  (Sb2S3). 

BISMUTH. 

Bismuth  is  usually  found  in  the  native  state,  though  it  is 
not  abundant  nor  widely  distributed.  The  oxide  (Bi2O3),  or 
bismite,  the  carbonate  ((BiO)2CO3.H2O),  or  bismutite,  and 
the  sulphide  (Bi2S3),  or  bismuthinite,  are  the  common  ores. 
The  world's  supply  comes  from  Saxony. 

Bismuth  is  prepared  from  the  native  metal  by  melting  it  on  an 
inclined  plate  and  allowing  it  to  drain  away  from  the  solid  impurities. 
Sometimes  the  sulphide  is  roasted,  and  the  resulting  oxide  is  reduced 
with  charcoal,  as  in  the  case  of  antimony. 

Bismuth  has  characteristic  properties.  It  is  a  grayish  white  metal 
with  a  reddish  tinge.  Like  antimony,  it  is  very  brittle.  It  does  not 
tarnish  in  dry  air,  but  it  grows  dull  in  moist  air ;  and  when  heated  in 
air  it  burns  with  a  bluish  flame,  forming  the  yellowish  oxide  (Bi2O3). 
Its  specific  gravity  is  about  9.9.  Hydrochloric  acid  does  not  readily 
attack  it,  but  nitric  acid  converts  it  into  a  nitrate,  and  hot  sulphuric  acid 
into  a  sulphate. 

Bismuth  melts  at  about  270°  C.  But  a  mixture  of  bismuth,  lead,  and 
tin  melts  at  a  low  temperature.  For  example,  Newton's  metal  melts  at 
95°  C.  and  Rose's  metal  at  100°  C. ;  while  Wood's  metal,  which  con- 
tains cadmium,  melts  at  only  66°  C.-yi0  C.  These  metallic  mixtures 
are  called  fusible  metals.  They  are  used  in  making  casts  of  wood 
cuts;  but  more  often  (i)  as  safety  plugs  in  steam  boilers  to  prevent 
explosions,  (2)  as  a  fuse  in  electrical  apparatus  to  prevent  a  short  cir- 
cuit, and  (3)  to  hold  in  place  fireproof  doors  and  the  valves  in  the 
automatic  sprinkling  apparatus  now  placed  in  large  buildings. 

Compounds  of  Bismuth.  —  Bismuth  forms  no  compound  with  hydro- 
gen. There  are  three  oxides.  Bismuth  trioxide  (Bi2O3)  is  yellowish, 


2j6  Descriptive  Chemistry. 

the  pentoxide  (Bi2O5)  is  orange  red,  and  the  dioxide  (Bi2(X)  is  black. 
Bismuth  trioxide  is  used  to  fix  the  gilding  on  porcelain.  The  trichloride 
(BiClt3)  is  formed  by  the  action  of  chlorine  upon  bismuth,  or  by  treat- 
ing bismuth  with  aqua  regia.  With  an  excess  of  water  the  trichloride 
forms  the  oxychloride  (BiOCl),  which  is  a  pearl-white  powder,  insoluble 
in  water.  The  formation  of  the  oxychloride  is  the  usual  test  for  bis- 
muth. Bismuth,  being  a  metal,  forms  hydroxides  (Bi(OH)3  and 
BiO.OH).  Normal  bismuth  nitrate  (Bi(NO.?);i),  treated  with  hot 
water,  forms  basic  bismuth  nitrate  (Bi(OH)2NCX  or  BiONO3).  The 
latter,  often  called  subnitrate  of  bismuth,  is  a  white  powder  used  as  a 
medicine  for  dyspepsia  and  as  a  cosmetic. 


EXERCISES. 

1 .  What  is  the  symbol  and  atomic  weight  of  phosphorus  ?     Give  a 
brief  history  of  this  element.     Why  is  it  so  named  ? 

2.  Discuss  the  occurrence  of  phosphorus. 

3.  Describe  the  manufacture  of  phosphorus  (a)  from  a  .phosphate 
and     sulphuric    acid,   and   (b)  by   the   electric   method.      How   is   it 
purified  ? 

4.  Summarize   the    properties   of  (a)    ordinary    phosphorus,   and 
(£)  red  phosphorus. 

5.  Describe  briefly  (a)  the  oxides  of  phosphorus,   (b}  orthophos- 
phoric  acid,  (c)  metaphosphoric  acid,  (d}  pyrophosphoric  acid,  (e)  phos- 
phine,  (/)  the  phosphorus  chlorides. 

6.  What  is  (a)  tricalcium  phosphate,  (b}  microcosmic  salt,  (c}  "  acid 
phosphate  "  ? 

7.  Describe  the  manufacture  of  (a)  sulphur  matches,  and  (b}  safety 
matches. 

8.  Discuss  the  relation  of  phosphorus  to  life. 

9.  What  is  a  fertilizer  ?     Name  three  natural  fertilizers.     Describe 
the  manufacture  of  artificial  fertilizer.     What  is  a  complete  fertilizer  ? 

10.  What  is  the  symbol  and  atomic  weight  of  arsenic  ? 

1 1 .  Name  several  ores  of  arsenic.     With  what  metals  is  arsenic  often 
associated  ? 

12.  Describe  the  preparation  and  state  the  properties  of  the  arsenic. 

13.  What  is  the  formula  of  arsenic  trioxide  ?     By  what  other  names 
is  it  known  ?     Summarize  its  properties.     For  what  is  it  used  ?     What 
is  the  antidote  for  arsenic  poisoning  ? 


Phosphorus,  Arsenic,  Antimony,  Bismuth.     277 

14.  What  is  (a)  Paris  green,  (6)  orpiment,  (V)  realgar  ?     For  what 
is  each  used  ? 

15.  Describe  Marsh's  test  for  arsenic. 

1 6.  What  is  the  symbol  and  atomic  weight  of  antimony  ? 

17.  In  what  forms  does  antimony  occur  and  where  is  it  found  ?     De- 
scribe its  preparation.     State  its  chief  properties. 

18.  What  is  tartar  emetic  ?     For  what  is  it  used  ? 

19.  Describe  the  test  for  antimony. 

20.  What  is  the  symbol  of  bismuth  ?     How  does  it  occur  and  where 
is  it  found  ?     Describe  its  preparation.     State  its  properties. 

21.  State  the  relation  of  bismuth  hydroxide  to  bismuth  subnitrate. 
Describe  the  latter. 

PROBLEMS. 

1.  Calculate  the  percentage  composition  of  (a)  sodium  phosphate 
(Na3PO4),  (£)  dihydrogen  phosphate  (H2NaPO4),  (V)  disodium  phos- 
phate (HNa2PO4),  O/)  microcosmic  salt  (HNaNH4PO4). 

2.  How  much  phosphorus  is   needed  to  remove  the  oxygen  from  a 
liter  of  air  ?     (Assume  (i)2P  +  5O  =  P2O5  and  (2)  air  is  20  per  cent 
oxygen.) 

3.  How  much  phosphorus  is  there  in  a  ton  (2000  Ib.)  of  bone  ash 
(Ca3(P04)2)? 

4.  If  a  skeleton  weighs  25  Ib.  and  contains  60  per  cent  calcium 
phosphate,  how  much  phosphorus  does  it  contain  ? 

5.  What  is  the  weight  of  a  cylindrical  stick  of  ordinary  phosphorus 
10  cm.  long  and  15  mm.  in  diameter  ?     (Suggestion.  —  What  is  the  spe- 
cific gravity  of  phosphorus  ?) 

6.  Calculate  the  percentage  composition  of  (a)  orpiment  (As3S3), 
($)  realgar  (As0S2),  (^)  white  arsenic  (As2O3). 

7.  What  is  the  weight  of  a  piece  of  antimony  25  cm.  long,  15  cm. 
wide,  and  2  mm.  thick  ? 


CHAPTER   XX. 
METALS. 

Introduction. — The  elements  studied  thus  far  are  chiefly 
non-metals.  Metals,  however,  have  been  mentioned,  and 
many  of  their  properties  have  been  discussed.  It  is  the 
purpose  of  the  present  chapter  to  review  these  properties 
and  prepare  the  way  for  a  fuller  treatment  of  the  metals. 

Metals  and  Non-metals.  —  Many  years  ago  the  chem- 
ical elements  were  divided  into  two  classes,  called  metals 
and  non-metals.  The  division  was  based  largely  on  the 
physical  properties  of  the  elements.  The  opaque,  lustrous, 
more  or  less  heavy,  hard,  ductile,  malleable,  tenacious 
solids  were  called  metals.  All  gases  and  the  solids  such 
as  carbon,  sulphur,  phosphorus,  and  iodine  were  called 
non-metals.  No  such  sharp  dividing  line,  however,  can 
be  drawn  between  metals  and  non-metals.  Some,  of 
course,  have  pronounced  properties,  like  the  non-metal 
sulphur  and  the  metal  iron.  These  are  typical.  But  a 
few  have  variable  properties.  Sometimes  they  act  as 
metals  and  at  other  times  as  non-metals.  Antimony  and 
arsenic  belong  to  this  border-line  class  ;  they  are  sometimes 
called  the  metalloids.  The  classification  into  metals  and 
non-metals  is  no  longer  accurate,  but  it  is  very  convenient. 
The  use  in  common  life  of  the  words  metallic  and  metal 
seldom  leads  to  confusion. 

Properties  of  Metals.  —  The  physical  properties  of 
metals  are  familiar,  though  variable  between  wide  limits. 

278 


Metals.  279 

All  have  a  metallic  luster,  i.e.  the  marked  property  of 
reflecting  light  from  their  polished  or  untarnished  surfaces. 

All  are  opaque  except  very  thin  films  of  gold.  The 
color  of  many  is  white,  though  the  tint  varies.  Thus 
silver,  sodium,  aluminium,  mercury,  magnesium,  iron,  and 
tin  are  nearly  pure  white,  and  bismuth  is  reddish  white. 
Copper  is  the  only  red  metal,  and  gold  the  only  yellow 
one,  which  is  an  element.  Most  metals  are  malleable  and 
ductile,  i.e.  they  may  be  hammered  or  rolled  into  sheets 
and  drawn  into  wire.  Gold,  copper,  silver,  iron,  platinum, 
and  aluminium  possess  both  these  properties  to  a  marked 
degree  ;  while  lead,  zinc,  and  tin  are  very  malleable  though 
not  so  ductile.  Antimony  and  bismuth  are  brittle.  The 
hardness  of  metals  varies.  At  the  ordinary  temperature 
mercury  is  a  liquid,  sodium  and  lead  can  be  cut  easily  with 
a  knife,  and  so  on  through  the  list  up  to  iridium,  which  is 
as  hard  as  steel.  In  specific  gravity,  which  was  once 
thought  must  very  high,  the  metals  range  between  lithium, 
which  has  the  specific  gravity  0.585,  and  osmium,  which  has 
the  specific  gravity  22.48.  Sodium  and  potassium  also  are 
lighter  than  water,  while  magnesium  has  the  specific  grav- 
ity 1.75,  and  aluminium  2.58.  Metals  are  good  conductors 
of  heat  and  electricity.  They  also  vary  in  this  property. 
Silver,  copper,  and  aluminium  are  the  best  conductors,  and 
have  therefore  many  practical  applications.  Bismuth  is 
the  poorest  conductor. 

The  distinctive  property  of  metals  is  not  physical,  but 
chemical.  Metals  form  oxides  which  combine  with  water 
to  produce  bases.  Metals  are  the  characteristic  elements 
of  bases.  On  the  other  hand,  non-metals  form  acid-pro- 
ducing compounds. 

Occurrence  of  Metals.  —  Only  a  few  metals  are  found 
free  in  the  earth's  crust,  and  these  are  seldom  pure.  Of 


280  Descriptive  Chemistry. 

the  six  metals  known  to  the  ancients,  —  gold,  copper,  silver, 
tin,  iron,  and  lead,  —  only  gold  and  copper  are  found  free. 
The  solid  elements  and  their  compounds  which  occur  in 
the  earth's  crust  are  called  minerals.  And  those  minerals 
from  which  metals  can  be  profitably  extracted  are  called 
ores.  The  most  abundant  classes  of  ores  are  oxides,  sul- 
phides, carbonates,  and  hydroxides.  Lead,  zinc,  mercury, 
and  silver  sulphides  are  abundant.  Besides  native  copper, 
the  sulphide  and  carbonate  are  found.  Iron  occurs  as 
oxide,  carbonate,  hydroxide,  and  sulphide.  Many  ores 
contain  arsenic.  Some  ores  are  very  complex. 

Preparation  of  Metals.  —  The  series  of  operations  by 
which  useful  metals  are  extracted  from  their  ores  is  called 
metallurgy.  It  includes  preliminary  treatment,  smelting, 
electrolysis,  refining,  and  other  operations  necessary  to 
change  the  ore  into  a  metal  ready  for  manufacture  into 
useful  articles.  The  object  of  the  preliminary  treat- 
ment is  to  prepare  the  ore  for  smelting  or  for  a  similar 
operation  by  which  the  metal  is  obtained  in  a  state 
adapted  for  further  purification  or  refining.  The  ore  as  it 
comes  from  the  mine  is  usually  mixed  with  earthy  matter 
or  rock  called  gangue.  This  impurity  is  removed  by  me- 
chanical or  chemical  processes,  sometimes  by  both.  The 
mechanical  process  illustrates  one  kind  of  preliminary  treat- 
ment. The  ore  is  first  crushed  in  a  stamp  mill.  This  is  a 
huge,  heavy  mortar  and  pestle.  The  pestle  or  stamp  falls 
repeatedly  upon  the  ore,  which  is  slowly  fed  into  the  mortar 
or  die.  A  current  of  water  (or  air)  forces  the  fine  particles 
out  of  the  mortar  through  a  sieve.  The  lighter  particles  of 
the  impurities  are  washed  away,  and  the  metallic  grains 
are  extracted  by  another  mechanical  operation,  though 
chemical  processes  are  frequently  employed,  especially 
with  inferior  ores.  This  separation  of  the  valuable  part 


Metals.  281 

of  the  ore  from  the  gangue,  and  reducing  it  to  a  smaller 
bulk  is  often  called  ore  dressing  or  concentration.  Copper 
is  extracted  from  the  Lake  Superior  ores  mainly  by  this 
method  of  preliminary  treatment. 

Gold  and  silver  ores  are  treated  this  way,  and  then  ex- 
tracted from  the  slime  by  mercury.  The  latter  operation 
is  called  amalgamation.  The  most  common  method  of 
extracting  metals  from  their  ores  is  by  smelting.  The 
process  varies  with  the  kind  and  composition  of  the  ore. 
Essentially,  it  consists  in  heating  a  mixture  of  the  ore  and 
coke  (or  coal)  in  a  furnace.  The  ores  used  must,  as  a  rule, 
be  oxides.  Sulphides,  hydroxides,  and  carbonates  are  first 
roasted  or  calcined  to  convert  them  into  oxides.  The 
essential  chemical  change  in  smelting  is  a  reduction  of  the 
oxide  by  the  carbon.  The  carbon  and  oxygen  unite  and 
pass  off  as  a  gas,  leaving  the  metal  to  run  out  at  the  bot- 
tom. Limestone,  or  a  similar  substance,  called  a  flux,  is 
added  to  the  mixture,  if  necessary,  to  facilitate  the  melting 
and  to  assist  in  removing  the  impurities  as  a  glassy  sub- 
stance, called  slag.  The  operation  is  conducted  in  differ- 
ent kinds  of  furnaces.  Iron,  for  example,  is  smelted  in  a 
huge  upright  furnace  called  a  blast  furnace  (Fig.  72), 
because  a  current  of  air  is  forced  through  the  melted  mass 
to  facilitate  the  fusion  and  chemical  changes.  In  such  a 
furnace  the  fuel  and  ore  are  in  direct  contact.  When  this 
is  undesirable,  the  reverberatory  furnace  is  used  (Fig.  54). 
As  the  figure  shows,  in  this  furnace  the  flame  is  reflected 
or  reverberated  upon  the  ore  under  treatment.  In  this 
kind  of  furnace  the  ore  may  be  oxidized  or  reduced  with- 
out coming  in  contact  with  the  fuel.  Some  ores  demand 
special  methods,  which  will  be  described  in  connection 
with  these  metals.  Electrolysis  is  used  to  extract  some 
metals,  especially  aluminium.  Other  metals,  notably 


282 


Descriptive  Chemistry. 


FIG.  54.  —  Reverberatory  furnace.    Tue  tire 


copper,  are  purified  by 
electrolysis.  A  few  met- 
als are  extracted  by  a 
wet  process.  That  is, 
the  ores  are  dissolved, 
and  the  metal  is  precipi- 
tated by  adding  some 
substance  or  by  elec- 


i  ivj.  5*j-»  -  AV<-  v^i  u\,i  aiwi  j    mi  lit  L\^...        i  nt;  me  -I  .  T*!.  *       X"          " 

burns  on  the  grate,   G,  and  the  long  flame     trolySlS.         1  huS,    interior 


which  passes  over  the  bridge,  E,  is  reflected  Pjold     Ores    are     dissolved 

down  by  the  sloping  roof  upon  the  contents  .  , 

of  the  furnace.    Gases  escape  through /.    The  by  treatment  With    potas- 

charge,  which  rests  upon  B,  does  not  come  sium      CV^nide      and     the 

in  contact  with  the  fuel,  but  is  oxidized  or  .  . 

reduced  by  the  flame.  gold   IS   then  precipitated 

by  zinc. 

Alloys  are  mixtures  or  compounds  of  two  or  more 
metals.  Some  fused  metals  mix  in  all  proportions,  while 
others  seem  to  form  definite  compounds.  The  properties 
of  alloys  vary  with  the  constituents  and  their  properties. 
Some  alloys,  especially  those  of  copper  and  of  lead,  have 
many  industrial  uses.  Alloys  in  which  mercury  is  a  con- 
stituent are  called  amalgams. 

EXERCISES. 

1.  Define  the  terms  metal  and  non-metal  as  they  are  ordinarily  used. 
Name  six  or  more  examples  of  each  class.     Define  and  illustrate  the 
term  metalloid.     Why  is  this  classification  inaccurate? 

2.  State   the   familiar  physical   properties   of  metals.     Define    (a) 
metallic  luster,  (b}  malleable,  (c)  ductile,  (d)  specific  gravity. 

3.  How  does  the  color  of  metals  differ  from  their  luster?     Name 
five  metals  which  are  white.     What  color  has  (a)  gold,  (6)  copper,  (c) 
zinc,  (a)  lead,  (e)  iron? 

4.  What  metals   are  brittle?     Malleable?     Soft?     Hard?     Heavy? 
Light?     What  metals  conduct  electricity  well? 

5.  What  is  the  distinctive  chemical  property  of  metals?     Of  nor>' 
metals?     Illustrate  your  answer. 


Metals.  283 

6.  What  metals  are  often  found  free  in  nature?     Define  and  illus- 
trate the  terms  (a)  mineral,  and  (b)  ore.     What  are  the  most  abundant 
classes  of  ores? 

7.  What  metals  occur  abundantly  as  (a)  sulphides,  (b)  oxides,  (V) 
carbonates  ? 

8.  Define  metallurgy.     WThat  general  operations  does  it  include? 

9.  What  is  the  object  of  the  preliminary  treatment  of  ores?     How 
is  it  accomplished  mechanically?     Define  (a)  gangue,  (<£)   concentra- 
tion, (c)  amalgamation.     What  metal  is  often  extracted  (a)  mechanic- 
ally, (b)  by  amalgamation  ? 

10.  Define  smelting.  What  fundamental  chemical  change  does  it 
usually  involve?  Define  and  illustrate  (a)  calcination,  ($)  flux,  (c) 
slag. 

n.  Describe  (a}  a  reverberatory  furnace,  and  (b)  a  blast  furnace. 
What  is  their  essential  difference  ? 

12.  What  is  the  wet  process  of  extracting  ores? 

13.  What  are  (a)  alloys,  (b)  amalgams? 

PROBLEMS. 

1.  What  is  the  specific  gravity  of  gold,  if  a  piece  weighs  4.676  gm. 
in  air,'  and  loses  0.244  gn».  when  weighed  in  water?     (Note.  —  Specific 
gravity  equals  the  weight  in  air  divided  by  the  loss  of  weight  in  water.) 

2.  A  piece  of  aluminium  weighs  150  gm.  in  air  and  75  gm.  in  water. 
What  is  its  specific  gravity? 

3.  A  piece  of  iron  weighs  292.8  gm.  in  air  and  255.3  gm.  in  water. 
What  is  its  specific  gravity? 

4.  A  piece  of  copper  weighing  50  gm.  in  air  lost  5.6  gm.  when 
weighed  in  water.     What  is  its  specific  gravity? 

5.  A  piece  of  lead  pipe  weighs  158.9  gm.  in  air  and  144.9  &m-  ^n 
water.     Calculate  the  specific  gravity. 


CHAPTER    XXL 
SODIUM,  POTASSIUM,  AND  LITHIUM. 

Introduction.  —  Sodium  and  potassium,  and  the  rare 
elements  lithium,  rubidium,  and  caesium,  form  a  natural 
group,  known  as  the  alkali  metals.  The  different  elements 
and  their  corresponding  compounds  resemble  each  other 
closely. 

Sodium  and  potassium  were  discovered  by  Sir  Humphry  Davy  in 
1807  by  the  electrolysis  of  their  hydroxides.  Bunsen,  by  means  of  the 
spectroscope,  discovered  lithium  in  1855,  caesium  in  1860,  and  rubidium 
in  1861. 

SODIUM. 

Occurrence.  —  Sodium  is  not  found  free.  Sodium  chlo- 
ride and  sodium  nitrate  are  the  most  abundant  compounds. 
Many  rocks,  plants,  and  mineral  waters  contain  combined 
sodium.  About  2.5  per  cent  of  the  earth's  crust  is  sodium. 

The  symbol  of  sodium,  Na,  is  from  the  Latin  word  natrium,  which 
in  turn  comes  from  the  Greek  word  natron,  an  old  name  of  sodium 
carbonate. 

Preparation.  —  Sodium  is  now  manufactured  on  a  large 
scale  by  the  electrolysis  of  fused  sodium  hydroxide.  This 
method  was  used  by  Davy  in  1807  to  isolate  sodium,  but 
its  commercial  success  was  only  recently  made  possible  by 
Castner.  Figure  55  is  a  sketch  of  the  apparatus  used. 
The  body  of  the  steel  cylinder,  S,  rests  within  a  heated 
flue.  Hence  the  sodium  hydroxide  is  solid  in  the  neck,  B, 
and  serves  to  protect  the  joint  made  by  the  iron  cathode, 

284 


SIR    HUMPHREY    DAVY 

1778-1829 

THE    FAMOUS    ENGLISH    CHEViST   WHOSE    BRILLIANT    DISCOVERIES    HAVE    NEVER    BEEN 
SURPASSED 


Sodium,   Potassium,  and  Lithium. 


285 


C,  and  the  crucible.  A,  A  is  the 
iron  anode.  A  collecting  pot, 
P,  dips  into  the  molten  caustic 
soda.  As  the  electrolysis  pro- 
ceeds, the  sodium  formed  at 
C  collects  in  P,  and  a  wire 
gauze,  G,  G,  keeps  it  from  mix- 
ing with  the  caustic  soda. 
The  sodium  is  ladled  out  at 
intervals  from  P.  The  hy- 
drogen, which  is  liberated, 
accumulates  also  in  P  and 
prevents  the  sodium  from  oxi- 
dizing. The  hydrogen  some- 
times escapes  and  explodes.  FIG.  55.  —  Apparatus  for  the  manu- 
facture of  sodium  by  the  electrolysis 

Sodium  was  formerly  manufactured    Of  sodium  hydroxide, 
by  two  methods,      (i)   Sodium   car- 
bonate and  carbon  heated  to  a  high  temperature  change  thus  — 


Na2C03 
Sodium  Carbonate 


2C     =      2Na 
Carbon        Sodium 


+  3  CO 

Carbon  Monoxide 


The  mixture  was  heated  in  iron  retorts,  and  the  sodium  vapor,  in  pass- 
ing through  a  flat  iron  receiver,  condensed  to  a  liquid,  which  was  col- 
lected under  paraffin  or  mineral  oil.  (2)  The  other  chemical  method, 
devised  by  Castner  in  1886,  consisted  essentially  in  heating  sodium 
hydroxide  with  a  mixture  of  iron  and  carbon.  Probably  iron  carbide 
was  the  essential  reducing  agent,  and  the  change  might  be  represented 
thus  — 

6NaOH      +      FeC2      =    2  Na    +   Fe  +    2  Na,CO,     +      3  H2 
Sodium  Hy-      Iron  Car-      Sodium      Iron      Sodium  Car-      Hydrogen 


droxide 


bide 


bonate 


Properties.  —  Sodium  is  a  silver-white  metal.  It  is  so 
soft  that  it  may  be  easily  molded  with  the  fingers  and 
cut  with  a  knife.  It  floats  upon  water,  since  its  specific 


286  Descriptive  Chemistry. 

gravity  is  only  0.98.  Heated  in  the  air,  it  melts  at  96°  C, 
and  at  a  higher  temperature  it  burns  with  a  brilliant  yellow 
flame,  forming  the  oxides  Na2O  and  Na2O2.  This  intense 
yellow  color  is  characteristic  of  sodium  and  is  the  usual 
test  for  the  element  (free  or  combined).  In  moist  air  the 
bright  surface  quickly  tarnishes,  and  sodium  as  usually 
seen  has  a  brownish  coating.  It  is,  therefore,  kept  under 
kerosene  or  a  liquid  free  from  water.  It  decomposes 
water  at  ordinary  temperatures,  liberating  hydrogen  and 
forming  sodium  hydroxide,  thus  — 

Na  +  H2O  -  NaOH  +  H 
Sodium  Water  Sodium  Hydroxide  Hydrogen 
If  held  in  one  place  upon  water  by  filter  paper,  enough 
heat  is  generated  to  set  fire  to  the  hydrogen,  which  burns 
with  a  yellow  flame,  owing  to  the  presence  of  volatilized 
sodium  (see  Interaction  of  Sodium  and  Water,  Chapter  V). 
If  melted  sodium  is  put  into  chlorine,  the  two  elements 
combine  with  a  brilliant  flame,  forming  sodium  chloride. 
It  was  in  this  way  that  Davy,  in  1810,  proved  that  com- 
mon salt  is  really  nothing  but  sodium  chloride.  It  combines 
directly  with  the  other  halogens. 

A  molecule  of  sodium  contains  only  one  atom. 

Sodium  is  used  in  the  laboratory  to  extract  water  from  alcohol  and 
ether  and  to  prepare  organic  compounds.  Large  quantities  are  con- 
sumed in  the  manufacture  of  sodium  peroxide  (NaaO2)  and  sodium 
cyanide  (NaCN).  Its  power  to  reduce  oxides  gives  it  limited  use  in 
preparing  certain  metals,  e.g.  magnesium. 

Sodium  Chloride,  NaCl,  is  the  most  important  compound 
of  sodium.  It  is  familiar  under  the  name  of  salt  or  com- 
mon salt.  The  presence  of  salt  in  the  ocean,  in  lakes  and 
springs,  and  in  the  soil  is  mentioned  in  the  oldest  histori- 
cal records.  It  is  one  of  the  most  abundant  substances. 
The  sources  of  salt  are  sea  water,  rock  salt,  and  brines. 


Sodium,  Potassium,  and  Lithium.  287 

Preparation  of  Salt.  —  Sea  water  contains  nearly  4  per  cent  of  salts, 
and  three  fourths  of  this  amount  is  sodium  chloride,  (i)  In  warm 
countries,  as  on  the  shores  of  the  Mediterranean  Sea,  shallow  ponds  of 
sea  water  near  the  shore  are  evaporated  by  exposure  to  the  sun  and 
wind,  and  the  salt  is  collected.  (2)  In  some  regions  sea  water  is  first 
concentrated  by  allowing  it  to  trickle  over  heaps  of  brush  and  then 
evaporated  to  crystallization  in  shallow  pans.  (3)  In  cold  countries, 
as  on  the  shores  of  the  White  Sea  in  Russia,  sea  water  is  allowed  to 
freeze  and  the  ice  is  removed.  The  ice  contains  no  salt,  so  the  opera- 
tion is  repeated  until  the  remaining  liquid  becomes  strong  enough  to 
evaporate  profitably  over  a  fire.  (4)  Deposits  of  salt  are  found  in 
many  parts  of  the  globe,  the  most  important  being  in  England,  Austria- 
Hungary,  and  Germany.  In  these  regions  and  some  parts  of  the 
United  States,  the  salt  is  mined  and  purified  like  other  minerals.  This 
variety  is  coarse  and  often  impure,  and  is  largely  used  in  curing  meat 
and  preserving  hides.  (5)  Most  of  the  salt  produced  in  the  United 
States  is  obtained  from  natural  or  artificial  brines,  i.e.  from  strong  solu- 
tions of  salt.  Artificial  brines  are  made  by  forcing  water  into  salt  de- 
posits. Brines  are  obtained  in  New  York,  Michigan,  Kansas,  Ohio, 
West  Virginia,  California,  Utah,  and  Louisiana.  They  are  evaporated 
in  vats  by  the  sun's  heat  or  by  heating  in  kettles  or  pans. 

All  these  methods  give  a  product  containing  as  impurities  salts  of 
sodium,  calcium,  and  magnesium,  which  are  largely  removed  by  further 
special  treatment.  The  dampness  of  salt  is  due  mainly  to  the  magne- 
sium chloride  it  contains  (see  Deliquescence,  Chapter  IV).  * 

Properties  and  Uses  of  Salt.  —  Salt  is  soluble  in  water, 
100  gm.  of  water  dissolving  about  36  gm.  of  salt  at  o°  C., 
and  40  gm.  at  100°  C.  It  crystallizes  in  cubes.  This  sub- 
stance is  an  essential  ingredient  of  the  food  of  man  and 
animals.  Besides  its  universal  domestic  use,  enormous 
quantities  are  consumed  in  the  preparation  of  many  so- 
dium compounds,  particularly  sodium  carbonate  (see  below), 
of  hydrochloric  acid  and  bleaching  powder.  In  1902  the 
United  States  produced  nearly  3,000,000  tons  of  salt,  and 
imported  over  200,000  tons.  This  is  about  the  average 
consumption. 


288  Descriptive  Chemistry. 

Sodium  Carbonate,  Na2CO3,  is  next  to  sodium  chloride 
in  importance.  Small  quantities  of  hydrated  sodium  car- 
bonates are  found  in  Egypt,  Russia,  and  in  California  and 
Nevada.  Formerly  it  was  obtained  from  the  ashes  of 
marine  plants,  but  sodium  chloride  is  now  the  source. 
The  manufacture  of  sodium  carbonate  is  one  of  the  most 
extensive  chemical  industries.  Two  processes  are  used, 
the  Leblanc  and  the  Solvay. 

The  Leblanc  Process  has  three  steps,  (i)  Sodium  chloride  is 
changed  into  sodium  sulphate  by  sulphuric  acid,  the  two  equations  for 
the  changes  being  — 

2NaCl      +       H2SO4      =       HNaSO4      +       HC1      +       NaCl 
Sodium  Sulphuric         Acid  Sodium   Hydrochloric     Sodium 

Chloride  Acid  Sulphate  Acid  Chloride 

HNaSO4         +         NaCl  Na,SO4         +         HC1 

Sodium  Sulphate 

This  operation  is  called  the  "salt  cake  process11;  the  impure  prod- 
uct, called  "  salt  cake,"  contains  about  95  per  cent  of  sodium  sulphate. 
The  hydrochloric  acid  is  a  by-product  (see  Hydrochloric  Acid) .  (2) 
The  sodium  sulphate  is  changed  into  sodium  carbonate  by  heating  the 
"salt  cake"  with  coal  and  limestone,  the  main  changes  being  repre- 
sented by  the  equations  — 

Na2S04  +  2  C  Na2S  +  2  CO, 

Sodium  Sulphate  Carbon          Sodium  Sulphide     Carbon  Dioxide 

Na2S          +          CaCO3  Na2CO3          +          CaS 

Sodium  Lime-  Sodium  Calcium 

Sulphide  stone  Carbonate  Sulphide 

This  operation  is  called  the  "black  ash -process."  The  product  is  a 
dark  brown  or  gray  porous  mass,  and  contains,  besides  37  to  45  per  cent 
of  sodium  carbonate,  considerable  calcium  sulphide  and  other  impuri- 
ties. The  calcium  sulphide  is  a  source  of  sulphur  (see  Sulphur).  (3) 
The  sodium  carbonate  is  rapidly  separated  from  the  insoluble  portions 
of  the  "  black  ash  "  by  agitation  with  a  small  amount  of  cool  water. 
The  solution  of  sodium  carbonate  thus  obtained  is  evaporated  to  crys- 


Sodium,  Potassium,  and   Lithium.  289 

tallization,  and  the  crude  crystals  are  ignited.  This  product  is  known 
as  soda  ash,  and  from  its  solution  in  waiter  are  obtained  soda  crystals 
or  sal  soda  (Na2CO3 .  10  H,O). 

The  Solvay  Process,  often  called  the  ammonia -soda  process,  con- 
sists in  saturating  a  cold  concentrated  solution  of  sodium  chloride  first 
with  ammonia  gas  and  then  with  carbon  dioxide  gas.  The  equation 
for  the  chemical  change  is—  .  it  0 

NaCl      +       NH3      +      CO2      =      HNaCO.,      +       NH4C1 
Sodium         Ammonia        Carbon        Acid  Sodium        Ammonium 
Chloride  Dioxide          Carbonate  Chloride 

The  acid  sodium  carbonate  is  nearly  insoluble  in  the  cold  ammonium 
chloride  solution,  and  therefore  separates.  It  is  changed,  by  heating, 
into  sodium  carbonate,  thus  — 

2  HNaCO,  Na2CO3          +          CO2         +        .  H2O 

Acid  Sodium  Sodium  Carbon  Water 

Carbonate  Carbonate  Dioxide 

The  liberated  carbon  dioxide  is  used  again,  and  from  the  ammonium 
chloride  the  ammonia  is  recovered  and  also  used. 

Properties  and  Uses  of  Sodium  Carbonate. —  Crystal- 
lized sodium  carbonate  (Na2CO3.  10  H2O)  is  often  called 
alkali  or  soda.  It  loses  water  in  the  air,  becoming  dull 
at  first  and  finally  falling  to  a  powder.  When  heated,  it 
melts  in  its  water  of  crystallization,  and  continued  heating 
changes  it  into  the  white  anhydrous  salt  (Na2CO3).  It  is 
readily  soluble  in  water,  and  the  solution,  which  is  strongly 
alkaline,  is  widely  used  as  a  cleansing  agent,  hence  the 
name  washing  soda. 

Enormous  quantities  of  sodium  carbonate  are  used  in 
the  glass  and  soap  industries,  and  in  preparing  sodium 
compounds. 

Sodium  Bicarbonate,  HNaCO3,  is  a  by-product  of  the 
Solvay  process,  and  it  may  also  be  prepared  by  treating 
crystallized  sodium  carbonate  with  carbon  dioxide  gas.  It 
is  a  white  powder,  less  soluble  in  water  than  the  normal 


290  Descriptive  Chemistry. 

carbonate.  When  heated  or  when  mixed  with  an  acid  or 
an  acid  salt,  sodium  bicarbonate  gives  up  carbon  dioxide. 
This  property  early  led  to  its  use  in  cooking,  and  gives  the 
names  cooking  soda,  baking  soda,  or  simply  soda. 

Sodium  bicarbonate  is  one  ingredient  of  baking  powder  and  of  the 
various  mixtures  (except  yeast)  used  to  raise  bread,  cake,  and  other 
food.  Since  cream  of  tartar  is  slightly  acid,- it  is  usually  used  to  liber- 
ate the  gas.  Sour  milk,  which  contains  lactic  acid,  is. sometimes  used 
in  place  of  cream  of  tartar.  When  pastry  is  raised  with  soda  and  cream 
of  tartar,  the  escaping  carbon  dioxide  puffs  up  the  dough.  Hence  bak- 
ing soda  is  often  called  saleratus  —  the  salt  which  aerates  (from  the 
Latin  words  sal,  salt,  and  aer,  air  or  gas).  Effervescing  powders,  such 
as  Seidlitz  (or  Rochelle)  and  soda  powders,  contain  sodium  bicarbon- 
ate in  one  paper  and  tartaric  acid  or  one  of  its  acid  salts  in  the  other. 
When  these  are  mixed  in  water,  carbon  dioxide  is  liberated.  Sodium 
bicarbonate  is  used  as  a  medicine  to  neutralize  an  acid  stomach.  For 
example,  the  "  soda  mints  "  sometimes  taken  for  this  purpose  are  mainly 
sodium  bicarbonate. 

Sodium  Hydroxide  or  Caustic  Soda,  NaOH,  is  a  white 
corrosive  solid.  It  absorbs  water  and  carbon  dioxide 
rapidly  from  the  air.  It  dissolves  readily  in  water,  with 
rise  of  temperature,  and  the  solution  is  strongly  alkaline. 
It  melts  easily,  and  is  often  cast  into  sticks  for  use  in  the 
laboratory.  Immense  quantities  are  used  in  making  hard 
soap,  paper,  and  dyestuffs ;  in  bleaching,  and  in  refining 
kerosene  oil. 

Sodium  hydroxide  is  usually  manufactured  by  treating  crude  sodium 
carbonate  with  calcium  hydroxide.  Lime  is  added  to  a  boiling,  dilute 
solution  of  soda  ash,  and  the  main  change  is  represented  thus  — 

Ca(OH)2     +      Na2CO3     =     2  NaOH     +     CaCO3 

Calcium  Sodium  Sodium  Calcium 

Hydroxide          Carbonate         Hydroxide       Carbonate 

The  solution  of  sodium  hydroxide  is  separated  from  the  insoluble  cal- 
cium carbonate,  and  concentrated  by  heating  in  iron  kettles  to  the  de- 


Sodium,  Potassium,  and  Lithium.  291 


sired  strength  or  until  the  mass  becomes  stiff.  Air  is  then  blown  in  or 
sodium  nitrate  added  to  oxidize  sulphides  to  sulphates.  After  standing 
several  hours  to  allow  other  impurities  to  settle,  the  caustic  soda  is  put 
into  iron  barrels  called  drums.  It  solidifies  on  cooling,  and  the  drums 
are  at  once  sealed  to  keep  out  the  air. 

Sodium  hydroxide  is  also  manufactured  on  a  large  scale 
at  Niagara  Falls,  New  York,  by  the  electrolysis  of  sodium 
chloride,  according  to  the  equation  — 


NaCl    + 

Sodium 
Chloride 


H2O    = 


NaOH  + 

Sodium 
Hydroxide 


Cl 

Chlorine 


+       H 

Hydrogen 


The  apparatus  is  shown  in  Figure  56.  The  carbon  anodes 
(A,  A)  pass  into  the  outer  compartments  which  contain 
brine,  and  the  iron  cathodes  into  the  middle  compartment 
which  contains  sodium  hydroxide  solution.  When  the  cur- 


FlG.  56.  —  Apparatus  for  the  manufacture  of  sodium  hydroxide  by  the  electrolysis 
of  sodium  chloride. 

rent  passes,  chlorine  is  evolved  at  the  anodes  and  flows  out 
through  pipes  (not  shown),  and  sodium  is  produced  on 
the  surface  of  the  mercury  (M )  which  covers  the  floor 
of  the  whole  apparatus.  The  sodium  forms  an  amalgam 


292  Descriptive  Chemistry. 

with  the  mercury,  and  by  rocking  the  apparatus  on  the 
device,  B,  B,  the  sodium  amalgam  flows  into  the  compart- 
ment, D,  where  the  sodium  is  liberated  by  the  action  of  the 
electric  current,  which  passes  between  the  cathode  and 
the  amalgam.  The  sodium  reacts  with  the  water  forming 
hydrogen,  which  passes  off  through  pipes  (not  shown)  and 
sodium  hydroxide,  which  flows  into  a  special  tank.  Both 
the  chlorine  and  sodium  hydroxide  are  nearly  pure.  The 
solution  of  caustic  soda  is  finally  treated,  if  necessary,  as  in 
the  older  process. 

Sodium  Sulphate,  Na2SO4,  is  one  of  the  products 
obtained  in  the  manufacture  of  sodium  carbonate  (see 
above). 

In  another  method,  sulphur  dioxide,  steam,  and  air  are  passed  into 
hot  sodium  chloride.  And  at  Stassfurt,  magnesium  sulphate  and  sodium 
chloride  are  allowed  to  interact  in  the  cold,  thus  — 

MgSO4      +    2NaCl      =      Na.SO4      +        MgCl2 

Magnesium         Sodium  Sodium  Magnesium 

Sulphate  Chloride  Sulphate  Chloride 

Sodium  sulphate  is  a  white  anhydrous  solid.  It  dissolves 
readily  in  water,  and  when  a  strong  solution  made  at  30°  C. 
is  cooled,  large  transparent  bitter  crystals  separate.  They 
have  the  formula  Na2SO4 .  ioH2O  and  are  called  Glau- 
ber's salt,  from  the  discoverer.  They  lose  water  when 
exposed  to  air,  and  the  salt  continues  to  effloresce  until  it 
becomes  an  anhydrous  powder.  The  crude  salt  is  used  in 
the  glass  and  dyeing  industries,  and  the  purified  salt  as  a 
medicine. 

Sodium  Nitrate,  NaNO3,  is  found  abundantly  in  Chili, 
and  is  often  called  Chili  saltpeter.  It  is  a  white  solid, 
which  becomes  moist  in  the  air.  Large  quantities  are  used 
as  a  fertilizer,  either  alone  or  mixed  with  compounds  of 


Sodium,   Potassium,  and  Lithium.  293 

potassium  and  of  phosphorus,  and  for  making  nitric  acid 
and  potassium  nitrate. 

The  natural  deposits  are  in  a  dry  region  near  the  coast  and  cover 
over  200,000  acres.  Chili  controls  the  industry,  and  exports  annually 
over  a  million  tons.  The  crude  salt,  which  looks  like  rock  salt,  is  puri- 
fied by  crystallization  into  a  product  containing  94-98  per  cent  of  the 
nitrate.  The  final  mother  liquor  is  a  source  of  iodine  (see  Iodine). 

Sodium  Dioxide  or  Peroxide,  Na.,Oa,  is  a  yellowish  solid.  It  is  used 
to  bleach  straw  and  delicate  fabrics.  With  water  it  liberates  oxygen, 
according  to  the  equation  — 

Na2O2          +     H20  O        +  2NaOH 

Sodium  Dioxide  Oxygen       Sodium  Hydroxide 

Miscellaneous.  —  Sodium  cyanide  (NaCN)  is  used  to  extract  gold 
from  poor  ores.  Sodium  monoxide  (Na2O)  is  a  grayish  solid.  The 
sodium  phosphates,  sodium  thiosulphate,  acid  sodium  sulphite,  sodium 
silicate,  and  sodium  tetraborate  or  borax  have  been  described. 

POTASSIUM. 

Occurrence.  —  This  metal  is  not  found  free,  but  its  com- 
pounds are  abundant.  The  minerals  mica  and  feldspar 
are  silicates  containing  potassium.  By  the  decay  of  these 
and  other  minerals,  potassium  compounds  find  their  way 
into  the  soil,  thence  into  plants  and  animals.  Potassium 
salts  are  found  in  wood  ashes,  in  suint,  —  the  oily  substance 
washed  from  sheep's  wool,  —  in  beet-sugar  residues,  and  in 
the  deposits  in  wine  casks.  Sea  water  and  mineral  waters 
contain  potassium  salts,  particularly  potassium  chloride 
and  potassium  sulphate.  Many  potassium  salts  are  found 
at  Stassfurt.  About  2.5  per  cent  of  the  earth's  crust  is 
potassium. 

The  Stassfurt  deposits  of  the  salts  of  potassium  and  other  metals 
are  near  Magdeburg,  Germany.  About  16  different  salts  make  up 
the  beds,  which  are  nearly  3000  feet  thick.  The  deposits  were  doubt- 
less formed  by  the  evaporation  of  sea  water,  though  the  different  simpler 


294  Descriptive  Chemistry. 

salts   interacted,  forming   complex  ones.     The   most   important  salts 

Kainite    ....  KC1,  MgSO4 .  3  H2O. 
Carnallite  .   .  .  KC1,  MgCl,, .  6  H2O. 
Polyhalite  .  .  .  K2SO4,  Mg~SO4,  2  CaSO4  .  2  H2O. 
Sylvite    ....  KC1.  . 

Picromerite  .  .  K2SO4,  MgSO4 .  6H2O. 

The  name  potassium  comes  from  the  word  potash.  The  symbol,  K, 
is  from  kalium,  the  Latin  equivalent  of  kali,  which  is  derived  from  an 
Arabic  term  for  an  alkaline  substance. 

Preparation.  —  Potassium  is  now  obtained  by  the  electrolysis  of 
potassium  hydroxide.  Formerly  it  was  manufactured,  like  sodium, 
by  heating  to  a  high  temperature  a  mixture  of  potassium  carbonate 
and  carbon  or  of  potassium  hydroxide  and  iron  carbide  (see  under 
Sodium). 

Properties.  —  Like  sodium,  potassium  is  a  soft,  silver- 
white  metal,  light  enough  to  float  upon  water  —  the  specific 
gravity  being  0.86.  Its  brilliant  luster  soon  disappears  in 
air,  owing  to  rapid  oxidation.  Potassium  as  ordinarily 
seen  is,  therefore,  covered  with  a  grayish  coating,  and,  like 
sodium,  must  be  kept  under  mineral  oil.  It  melts  at  62.5° 
C,  and  at  a  higher  temperature  burns  with  a  violet-colored 
flame.  This  color  is  characteristic  of  burning  potassium, 
arid  is  a  test  for  the  metal  and  its  compounds.  Like 
sodium,  it  decomposes  water  at  ordinary  temperatures, 
though  more  energetically.  The  heat  evolved  immediately 
ignites  the  hydrogen,  and  the  melted  potassium  surrounded 
by  a  violet  flame  dashes  to  and  fro  upon  the  cold  water. 
The  main  reaction  corresponds  to  the  equation  — 

K        -f  H20     =  KOH  +         H 

Potassium         Water  Potassium  Hydroxide         Hydrogen 

Potassium  combines  with  the  halogens  and  other  ele- 
ments more  vigorously  than  sodium,  and  forms  analogous 
compounds. 


Sodium,   Potassium,  and   Lithium.  295 

Potassium  Chloride,  KC1,  is  found  native  in  the  Stass- 
furt  deposits.  It  is  also  obtained  in  large  quantities  by 
decomposing  carnallite  and  crystallizing  the  potassium 
chloride  from  the  more  soluble  magnesium  chloride.  It  is 
a  white  solid  which  crystallizes  in  cubes  and  otherwise 
resembles  sodium  chloride.  It  is  used  chiefly  to  prepare 
other  potassium  salts,  especially  the  nitrate  and  chlorate. 

Potassium  bromide  and  potassium  iodide  have  been  described  (see 
Chapter  XVI). 

Potassium  Nitrate,  KNO3,  is  also  called  niter  and  salt- 
peter. It  is  formed  in  the  soil  of  many  warm  countries 
by  the  decomposition  of  nitrogenous  organic  matter  (see 
Nitrification). 

It  is  now  made  by  mixing  hot,  concentrated  solutions  of  native  so- 
dium nitrate  and  potassium  chloride,  which  interact  thus  — 

NaNO3  +  KC1  KN03  +  NaCl 

Sodium  Potassium  Potassium  Sodium 

Nitrate  Chloride  Nitrate  Chloride 

The  sodium  chloride,  being  less  soluble,  separates,  and  is  removed.  By 
evaporation,  small  crystals  of  potassium  nitrate,  called  "  niter  meal,11  are 
obtained,  and  further  purified  by  recrystallization. 

Potassium  nitrate  is  a  white  solid.  It  dissolves  easily  in 
cold  water  with  a  fall  of  temperature,  and  very  freely  in  hot 
water,  but  it  is  not  hygroscopic.  It  is  crystalline,  but  con- 
tains no  water  of  crystallization.  The  taste  is  salty  and 
cooling.  It  melts  at  339°  C.,  and  further  heating  changes 
it  into  potassium  nitrite  (KNO2)  and  oxygen.  At  a  high 
temperature,  potassium  nitrate  gives  up  oxygen  readily, 
especially  to  charcoal,  sulphur,  and  organic  matter.  This 
oxidizing  power  leads  to  its  extensive  use  in  making  gun- 
powder, fireworks,  matches,  explosives,  and  in  many  chemi- 
cal operations. 


296  Descriptive  Chemistry. 

Gunpowder  is  a  mixture  of  potassium  nitrate,  charcoal,  and  sulphur. 
The  ingredients  are  first  purified,  pulverized,  and  thoroughly  mixed. 
This  mixture  is  pressed,  while  damp,  into  a  thin  sheet ;  and  the  "  press 
cake"  thus  formed  is  broken  into  small  grains,  which  are  sorted  by 
sieves.  The  grains  are  then  smoothed  or  ''glazed"  by  rolling  them  in 
a  barrel,  again  sifted,  arid  finally  dried  at  a  low  temperature.  The  pro- 
portions differ  with  the  use  of  the  powder.  The  United  States  army 
standard  black  powder  contains  75  per  cent  of  potassium  nitrate,  15  of 
charcoal,  and  10  of  sulphur.  When  gunpowder  burns  in  a  closed  space, 
a  large  volume  of  gas  is  suddenly  formed.  So  enormously  is  this  gas 
expanded  by  the  heat  that  it  would  fill  several  hundred  times  the  space 
taken  by  the  powder  itself.  The  pressure  exerted  by  this  expanding  gas 
is  many  tons.  It  is  this  pressure  which  forces  the  ball  from  a  cannon 
and  tears  a  rock  to  pieces.  The  chemical  changes  attending  the  explo- 
sion of  gunpowder  in  a  closed  space  are  complex,  as  may  be  seen  by  the 
following  (approximate)  equation  :  — 

8  KNO3  +  90  +  38  =  2  K,C03  +  K2SO4  -f  K2S.,  +  7  CO2  +  4  N2 

Probably  secondary  reactions  produce  other  gases  besides  carbon  diox- 
ide and  nitrogen. 

Potassium  Chlorate,  KC1O3,  is  a  white,  crystallized,  lus- 
trous solid.  It  tastes  like  potassium  nitrate.  It  melts  at 
334°  C.,  and  at  a  high  temperature  decomposes  into  oxygen 
and  potassium  chloride  as  final  products,  thus  — 

KC1O3  KC1  +       O3 

Potassium  Chlorate  Potassium  Chloride  Oxygen 

It  is  used  to  prepare  oxygen,  and  in  the  manufacture  of 
matches  and  fireworks.  In  the  form  of  "  chlorate  of  potash 
tablets  "  it  is  used  as  a  remedy  for  sore  throat. 

Potassium  chlorate  is  manufactured  by  passing  chlorine  into  calcium 
hydroxide  (milk  of  lime)  and  adding  potassium  chloride  to  the  mixture. 
The  simplest  equations  for  the  complex  changes  may  be  written  thus  :  — 

(i)  6  Ca(OH)2  +  6  C12  =   Ca(ClO3)2         +          5  CaCl2      +     6  H,O 
Calcium  Calcium  Calcium 

Hydroxide  Chlorate  Chloride 


Sodium,  Potassium,  and  Lithium.  297 

(2)  3  Ca(ClO)2  Ca(C103)2  +         2  CaCl2 
Calcium  Hypochlorite             Calcium  Chlorate 

(3)  Ca(ClO3)2       +       2  KC1  2  KC1O3  +         CaCl2 

Potassium  Chlorate 

The  salt  is  also  made  by  the  electrolysis  of  a  hot  solution  of  potassium 
chloride,  though  it  has  been  found  more  satisfactory  to  first  prepare 
sodium  chlorate  and  convert  this  salt  into  potassium  chlorate  by  po- 
tassium chloride. 

Potassium  Carbonate,  K2CO3,  is  a  white  powder.  It 
deliquesces  in  the  air,  is  very  soluble  in  water,  and  the 
solution  has  a  strong  alkaline  reaction.  It  was  formerly 
obtained  by  treating  wood  ashes  with  water,  and  evaporating 
the  solution  to  dryness.  The  crude  salt  thus  obtained  has 
long  been  called  potash,  and  a  purer  product  is  known 
as  pearlash.  (The  term  potash  is  sometimes  applied  to 
potassium  oxide,  K2O.)  It  is  used  extensively  in  the  manu- 
facture of  hard  glass,  soft  soap,  caustic  potash,  and  other 
potassium  compounds. 

Potassium  carbonate  is  obtained  from  suint  by  igniting  the  greasy 
mass  and  extracting  the  potassium  carbonate  with  water.  Beet-sugar 
residues  also  furnish  potassium  carbonate.  After  the  sugar  has  been 
obtained  from  the  beet  sirup,  the  molasses  is  changed  by  fermentation 
into  alcohol,  which  is  distilled  off;  the  liquid  residue  is  evaporated  to 
dryness  and  ignited,  and  the  potassium  carbonate  extracted  with  water. 
Pure  potassium  carbonate  is  prepared  by  igniting  cream  of  tartar  made 
from  the  deposits  in  wine  casks.  All  these  sources  emphasize  the  inti- 
mate relation  of  potassium  compounds  to  vegetable  and  animal  life. 
The  bulk  of  the  potassium  carbonate  is  now  made  from  potassium  sul- 
phate or  from  the  chloride  by  the  Leblanc  process,  owing  to  the  abun- 
dance of  crude  potassium  salts  at  Stassfurt. 

Potassium  Hydroxide  or  Caustic  Potash,  KOH,  is  a 

white  brittle  solid,  resembling  caustic  soda.  It  absorbs 
water  and  carbon  dioxide  very  readily ;  and  if  exposed 
to  the  air,  soon  becomes  a  thick  solution  of  potassium 


298  Descriptive  Chemistry. 

carbonate.  Like  sodium  hydroxide,  it  dissolves  in  water 
with  evolution  of  heat,  forming  a  strongly  alkaline  caustic 
solution.  It  is  one  of  the  strongest  bases,  even  glass  and 
porcelain  being  corroded  by  it.  Besides  its  use  in  the  labo- 
ratory, large  quantities  are  consumed  in  making  soft  soap. 

Potassium  hydroxide  is  made  and  purified  in  the  same  way  as  sodium 
hydroxide,  viz.  by  adding  lime  or  milk  of  lime  to  a  boiling  dilute  solution 
of  potassium  carbonate,  the  equation  for  the  change  being :  — 

Ca(OH)2  -f  K,CO3  2  KOH         +         CaCO3 

Milk  of  Potassium  Potassium  Calcium 

Lime  Carbonate  Hydroxide  Carbonate 

It  is  also  made  by  the  electrolysis  of  a  solution  of  potassium  chloride. 

Miscellaneous.  —  Potassium  Cyanide  (KCN)  is  a  white  solid,  very 
poisonous,  very  soluble  in  water,  and  having  an  odor  like  bitter  almonds 
(see  Cyanogen,  Chapter  XIV) .  Potassium  Sulphate  (K2SO4)  is  manu- 
factured from  kainite,  and  is  largely  used  as  a  fertilizer  and  in  making 
potassium  carbonate. 

Relation  of  Potassium  to  Life.  —  Potassium,  like  nitro- 
gen and  phosphorus,  is  essential  to  the  life  of  plants  and 
animals.  The  ash  of  many  common  grains,  vegetables,  and 
fruits  contains  potassium  as  the  carbonate.  Potassium  salts 
are  supposed  to  assist  in  the  formation  of  starch,  just  as 
phosphorus  is  indispensable  to  the  transformation  of  nitro- 
gen compounds.  Potassium  salts  taken  from  the  soil  by 
plants  must  be  returned  if  the  soil  is  to  be  productive. 
Sometimes  crude  kainite  is  used  extensively  as  a  fertilizer  ; 
but  wood  ashes,  or  the  sulphate  and  chloride,  are  often 
used  to  supply  potassium  salts. 

Lithium,  Li,  is  a  silver-white  metal  and  has  the  specific  gravity  of 
only  0.59.  It  is  the  lightest  of  the  metallic  elements.  Its  compounds 
are  widely  distributed  in  small  quantities  in  minerals,  mineral  waters, 
and  plants.  Lithia  water  and  citrate  of  lithium  are  often  prescribed  as 
a  remedy  for  diseases  of  the  kidneys.  Lithium  compounds  color  the 
Bunsen  flame  bright  red  —  a  delicate  test  for  the  metal. 


Sodium,  Potassium,  and  Lithium.  299 

Rubidium  and  Caesium,  Rb  and  Cs,  have  properties  and  form  com- 
pounds analogous  to  those  of  potassium. 


EXERCISES. 

1.  Name  the  alkali  metals.     What  is  the  symbol  of  each  ?     When 
and  by  whom  was  each  discovered  ? 

2.  What  are  the  important  compounds  of  sodium  ?     What  per  cent 
of  the  earth's  crust  is  sodium  ? 

3.  Describe  the  manufacture  of  sodium  by  electrolysis.     Describe 
the  older  methods  of  manufacture. 

4.  Summarize  (#)  the  physical  properties,  and  (<£)  the  chemical 
properties  of  sodium.     How  is  it  usually  kept  ?     For  what  is  it  used  ? 

5.  Discuss  the  interaction  of  sodium  and  water  (see  Chapter  V). 

6.  Give  the  chemical  name  and  formula  of  common  salt     Where  is 
it  found  ? 

7.  Describe  the  different  methods  of  preparing  salt.     State  (#)  the 
properties,  and  (b)  the  uses  of  salt. 

8.  Discuss  the  manufacture  of  sodium  carbonate  by  (a)  the  Le- 
blanc  process,     (b)  By  the  Solvay  process. 

9.  What  is  (a)  soda,  (b)  soda  ash,  (V)  salt  cake,  (</)  soda  crystals, 
(V)  sal  soda,  (/*)  washing  soda,  (g)  "alkali"  ? 

10.    State  the  properties  and  uses  of  sodium  carbonate. 
n.   Describe  the  preparation,  and  state  (a)  the  properties,  and  (£) 
the  uses  of  sodium  bicarbonate. 

12.  What  is  (a)  acid  sodium  carbonate,  (b)  saleratus,  (c}  baking 
powder,  (d)  baking  soda,  (e)  caustic  soda  ? 

13.  State  the  properties  and  uses  of  sodium  hydroxide. 

14.  Describe  the  manufacture  of  sodium  hydroxide  (a)  from  lime 
and  sodium  carbonate,  and  (b)  by  electrolysis  of  sodium  chloride. 

15.  How  is  sodium  sulphate   manufactured?     State  its  properties 
and  uses. 

1 6.  Where  is  sodium  nitrate  found  ?    State  its  properties  and  uses. 

17.  Review  briefly  (a)  sodium   thiosulphate,  (b}  water  glass,   (c) 
borax. 

1 8.  What  is  a  simple  test  for  (a)  sodium,  and  (£)  potassium  ? 

19.  Give  the  formula  of  (#)  sodium  carbonate,  (<£)  sodium  chloride, 
(c)  sodium  sulphate,  (d)  sodium  hydroxide,  (e)  sodium  bicarbonate, 
(/)  Glauber's  salt,  (g)  sodium  nitrate. 


300  Descriptive  Chemistry. 

20.  Discuss  the  occurrence  of  potassium  compounds. 

21.  Discuss  the  Stassfurt  deposits.    . 

22.  How  is  potassium  prepared  ?     State  (a)  its  physical  properties, 
and  (£)  its  chemical  properties. 

23.  Describe  the  interaction  of  potassium  and  water. 

24.  Describe  the  preparation,  and  state  the  properties  and  uses  of 
(#)  potassium  chloride,  and  (<£)  potassium  nitrate. 

25.  Compare  potassium  nitrate  and  potassium  nitrite. 

26.  Describe  the  manufacture  of  gunpowder.     Upon  what  does  its 
use  depend  ? 

27.  State  the  properties  and  uses  of  (#)   potassium  chlorate,  ($) 
potassium  carbonate,  (c)  potassium  hydroxide. 

28.  Describe  the  manufacture  of  (a)  potassium  chlorate,  (£)  potas- 
sium carbonate,  (c)  potassium  hydroxide. 

29.  What  is  (a)  potash,  (£)  pearlash,  (c}  chlorate  of  potash  ? 

30.  Discuss  the  relation  of  potassium  to  life. 

31.  State  the  derivation  of  the  names  {a)  sodium,  and  (b)  potassium. 

32.  What  is  (a)  niter,  (£)  saltpeter,  (c)  Chili  saltpeter  ? 

33.  What  is  the  formula  of  the  following  compounds  of  potassium : 
(#)  hydroxide,  (b)  carbonate,  (c)  nitrate,  (</)  nitrite,  (e)  sulphate,  (/) 
chlorate,  (g)  cyanide  ? 

34.  Describe  lithium.     For  what  are  its  compounds  used  ? 


PROBLEMS. 

1.  How  much  potassium  carbonate  is  necessary  to  prepare  a  kilo- 
gram of  potassium  hydroxide?     (Assume  K9CO3  +  Ca(OH).,=  2  KOH 
+  CaC08.) 

2.  What  per  cent  of  Glauber's   salt,  Na2SO4 .  ioH2O,   is   sodium 
sulphate  ? 

3.  A  gram  of  gunpowder  produced  300  cc.  of  gas  at  o°  C.     What 
would  be  the  volume  at  2300°  C.  ? 

4.  How  much  sodium  will  2  kg.  of  sodium  carbonate  yield,  if  heated 
with  carbon  ?     (Assume  Na2CO3  +  C2  =  Na2  +  3  CO.) 

5.  What  is  the  per  cent  of  sodium  in  (a)  NaOH/  (£)  Na2SO4,  (c} 
NaCl,  (d)  HNaSO4  ? 

6.  What  is  the  per  cent  of  potassium  in  (#)  potassium  bromide 
(KBr),  (£)  potassium  nitrate  (KNO3),  (c)  potassium  iodide  (KI)  ? 


CHAPTER  XXII. 
COPPER  —  SILVER  —  GOLD. 

Introduction. — r These  metals  are  related,  but  they  do 
not  form  a  group  having  such  marked  family  character- 
istics as  the  alkali  metals.  The  metals,  as  well  as  their 
alloys  and  compounds,  have  many  domestic  and  commer- 
cial uses. 

COPPER. 

Copper  has  been  known  for  ages.  Domestic  utensils 
and  weapons  of  war  containing  copper  were  used  before 
similar  objects  of  "iron.  The  Romans  obtained  copper 
from  the  island  of  Cyprus/  They  called  it  cuprium  aes 
(i.e.  Cyprian  brass),  which  finally  became  simply  cuprum. 
From  cuprum  we  obtain  the  symbol  Cu  and  the  terms  cu- 
prous and  cupric. 

Occurrence  of  Copper.  —  Copper,  both  free  and  com- 
bined, is  an  abundant  element.  Single  masses  of  native 
or  metallic  copper  weighing  many  tons  are  found  in  Michi- 
gan mines  on  the  shores  of  Lake  Superior.  The  most 
valuable  ores  of  copper  are  copper  sulphide  (chalcocite, 
copper  glance,  Cu2S),  copper  oxide  (cuprite,  ruby  ore, 
Cu2O),  the  copper-rron  sulphides  (copper  pyrites,  chal- 
copyrite,  CuFeS2,  and  bornite,  Cu3FeS3),  and  the  conir 
plex  carbonates  (malachite,  CuCO3Cu(OH)2,  and  azurite, 
2  CuCOg  .  Cu(OH)2). 

Native  copper  conies  chiefly  from  Michigan  (Fig.  71),  the  copper- 
iron  sulphide  ores  from  Montana,  and  the  carbonates  from  Arizona. 

301 


302  Descriptive  Chemistry. 

The  United  States  produced  about  300,000  tons  of  copper  in  1902, 
which  was  more  than  half  of  the  world's  supply.  Of  this  amount  Mon- 
tana furnished  about  38  per  cent,  Michigan  26  per  cent,  and  Arizona  22 
per  cent.  The  annual  output  has  steadily  increased  since  1896. 

Metallurgy  of  Copper.  —  Copper  is  extracted  from  its 
ores  by  processes  which  vary  with  the  composition  of  the 
ore.  (i)  Native  copper  ore  is  first  crushed,  then  washed 
to  remove  impurities,  and  the  concentrated  product  finally 
smelted  and  refined  by  a  single  fusion.  (2)  The  carbon- 
ates and  oxides  are  reduced  by  roasting  them  with  coke  in 
blast  furnaces.  The  general  chemical  change  may  be  rep- 
resented thus  — 

Cu2O        +       C      =    2Cu    +  CO 

Copper  Oxide         Carbon          Copper        Carbon  Monoxide 

(3)  The  smelting  of  copper-iron  sulphides  is  complicated. 
The  ore  is  crushed  and  washed,  and  then  roasted  in  a  fur- 
nace. This  operation  removes  the  adhering  rock  and 
changes  much  of  the  sulphide  into  an  oxide.  The  roasted 
mass  is  then  melted  with  coal  and  sand  in  a  shaft  or  a 
reverberatory  furnace,  whereby  the  iron  is  largely  changed 
into  a  fusible  silicate,  which  runs  off  as  a  part  of  the  slag. 
The  remaining  "matte,"  as  it  is  called,  contains  from  50 
to  65  per  cent  of  copper,  besides  some  iron,  sulphur,  and 
arsenic.  It  is  roasted  and  melted  until  all  the  iron  and 
arsenic  are  removed  and  mainly  copper  sulphide  remains. 
This  is  finally  roasted  to  convert  it  partly  into  an  oxide, 
and  the  mixture  of  sulphide  and  oxide  is  again  melted  ;  the 
sulphur  passes  off  as  sulphur  dioxide,  and  the  copper  is 
left  behind.  The  equation  for  this  final  change  is  — 

2  CuO        +        Cu2S       =       4  Cu       -h       SO2 
Copper  Oxide        Copper  Sulphide          Copper        Sulphur  Dioxide 


Copper  —  Silver  —  Gold. 


303 


1 

0 

1 

o 

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0 

± 

m 

^ 

A 

A 

A 

A 

s 

r 

3 

r— 

~^ 

-- 

-~- 

- 

; 

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! 

1 

j£ 

£p; 

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^£ 

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- 

; 

Sometimes  the  sulphur  and  arsenic  are  removed  by  forcing 
hot  air  through  the  molten  sulphide. 

Purification  of  Copper.  —  The  crude  copper  from  most 
ores  contains  about  98  per  cent  of  copper.  Such  impure 
copper  is  best  purified  by  electrolysis,  and  is  called  electro- 
lytic copper.  Thick  plates  of  the  impure  copper  are 
attached  as  anodes  to  the 
positive  electrode  of  a 
powerful  battery  or  dy- 
namo and  suspended  in 
a  solution  of  copper  sul- 
phate and  sulphuric  acid. 
Sheets  of  pure  copper  are 

attached     as     Cathodes     tO          FIG.  57.  — Apparatus  for  the  preparation 

the  negative  electrode  and,  *££&  *  S±hodet  *  *  *" 
of  course,  dip  into  the 

solution  (Fig.  57).  When  the  current  passes,  the  crude 
copper  anodes  dissolve,  pure  copper  is  deposited  upon  the 
cathodes,  and  the  impurities  either  remain  in  solution  or 
fall  to  the  bottom  of  the  tank  as  mud.  From  this  mud, 
gold  and  silver  are  extracted  in  appreciable  quantities. 
Electrolytic  copper  is  very  pure. 

Properties  of  Copper.  —  Copper  is  a  bright  metal,  dis- 
tinguishable from  all  others  by  its  peculiar  reddish  color. 
It  is  flexible,  hard,  and  tough ;  it  can  be  drawn  out  into 
wire  and  rolled  into  very  thin  sheets.  Its  specific  grav- 
ity is  8.9.  Next  to  silver,  copper  is  the  best  conductor  of 
heat  and  electricity.  Exposed  to  dry  air,  it  turns  dull,  and 
in  moist  air  it  gradually  becomes  coated  with  a  greenish 
copper  carbonate.  Heated  in  the  air,  it  is  changed  into 
the  black  copper  oxide,  and  at  a  high  temperature  it  colors 
a  flame  emerald  green.  With  nitric  acid  it  forms  copper 


304  Descriptive  Chemistry. 

nitrate  and  oxides  of  nitrogen  (see  Oxides  of  Nitrogen); 
with  hot  sulphuric  acid  it  yields  copper  sulphate  and  sul- 
phur dioxide  (see  Sulphur  Dioxide).  Hydrochloric  acid 
has  little  effect  upon  it.  Copper  replaces  some  metals  if 
suspended  in  solutions  of  their  compounds,  e.g.  a  clean 
copper  wire  soon  becomes  coated  with  mercury  if  placed 
in  a  solution  of  any  mercury  compound ;  on  the  other 
hand,  metals  like  iron,  zinc,  and  magnesium  remove  cop- 
per from  its  solution,  e.g.  a  nail  or  knife  blade  soon  becomes 
coated  with  copper  if  dipped  into  a  solution  of  any  copper 
compound.  Scrap  iron  is  often  used  to  precipitate  copper 
on  a  large  scale. 

Test  for  Copper. —  (i)  The  reddish  color,  peculiar  "coppery"  taste, 
and  green  color  given  to  a  flame  serve  to  identify  metallic  copper. 
(2)  An  excess  of  ammonium  hydroxide  added  to«the  solution  of  a  cop- 
per compound  produces  a  beautiful  blue  solution.  (3)  A  few  drops 
of  acetic  acid  and  of  potassium  ferrocyanide  solution  added  to  a  dilute 
solution  of  a  copper  compound  produce  a  brown  precipitate  of  copper 
ferrocyanide.  These  tests  are  characteristic  and  decisive. 

Uses  of  Copper.  —  Next  to  iron,  copper  is  the  most  use- 
ful metal.  Enormous  quantities  of  wire  are  used  in  operat- 
ing the  telegraph,  cable,  telephone,  electric  railway,  and 
electric  light.  Sheet  copper  is  made  into  household  utensils 
boilers,  and  stills.  Copper  bolts,  nails,  and  rivets  are  used 
on  ships,  because  the  rust  does  not  destroy  wood  as  iron  rust 
does.  All  nations  use  copper  as  the  chief  ingredient  of 
small  coins.  Electrical  apparatus  utilizes  much  copper. 
Maps,  etchings,  and  some  kinds  of  engravings  are  printed 
from  copper  plates ;  calico  is  printed  from  a  copper  cyl- 
inder upon  which  the  design  is  engraved.  Books  are 
printed  and  illustrated  from  an  electrotype,  made  by  de- 
positing a  film  of  copper  upon  an  impression  of  the  type 
or  design  in  wax  or  plaster  of  Paris.  In  a  similar  way 


Copper  —  Silver  —  Gold.  305 

many  objects  are  copper  plated  (see  Chapter  X).     Copper 
is  an  essential  constituent  of  many  alloys. 

Alloys  of  Copper  are  important.  Brass  is  a  bright  yel- 
low alloy  containing  63  to  72  per  cent  of  copper,  the  re- 
mainder being  zinc.  It  is  made  by  melting  these  metals 
together.  It  can  be  drawn  into  wire,  hammered  into  any 
shape,  and  turned  in  a  lathe.  It  is  harder  than  copper, 
and  on  account  of  its  durability  and  elasticity  has  many  uses 
for  which  copper  is  not  suited.  Pinchbeck,  Muntz  metal, 
Bath  metal,  Dutch  metal  (leaf  or  "gold"),  are  varieties  of 
brass.  Muntz  metal  is  now  used  in  place  of  sheet  copper, 
as  sheathing  for  the  bottoms  of  ships,  because  it  rusts  very 
slowly.  Typical  bronze  contains  different  proportions  of 
copper,  zinc,  and  tin.  Some  antique  bronzes  contain  lead 
or  iron.  The  per  cent  of  copper  is  70  to  95,  of  zinc  I  to 
25,  of  tin  i  to  1 8.  The  proportions  in  the  British  bronze 
coinage  are  copper  95,  zinc  i,  tin  4.  On  account  of  its 
beautiful  color  and  extreme  durability,  bronze  is  used  for 
statues,  memorial  tablets,  coins,  and  medals.  The  ancients 
made  it  into  weapons  of  war  and  household  utensils.  Can- 
non were  formerly  made  of  bronze,  but  for  this  purpose 
steel  is  now  used.  Phosphor  bronze  contains  a  small  per 
cent  of  phosphorus  and  of  lead.  It  is  tougher  than  ordi- 
nary bronze,  and  is  used  to  make  steamship  propellers  and 
parts  of  machines.  Silicon  bronze  is  copper  with  traces  of 
iron  and  silicon,  and  is  used  for  telegraph  and  telephone 
wires.  Aluminium  bronze  contains  90  per  cent  copper 
and  10  per  cent  aluminium.  It  is  a  hard,  yellow,  elastic 
alloy,  and  is  used  in  constructing  hulls  of  yachts  ;  its  light- 
ness, strength,  and  resistance  to  chemicals  adapt  it  to  many 
other  uses. 

Gun  metal  is  about  90  per  cent  copper  and  10  per  cent  zinc  ;  it  was 
formerly  used  in  making  cannon,  and  is  now  used  to  some  extent  in 


306  Descriptive  Chemistry. 

making  firearms.  Bell  metal  contains  about  75  per  cent  copper  and 
25  per  cent  zinc.  Speculum  metal  contains  about  70  per  cent  copper, 
30  per  cent  tin,  and  traces  of  zinc,  nickel,  and  iron  ;  it  takes  a  brilliant 
polish,  and  is  used  in  optical  instruments,  especially  telescopes,  to  re- 
flect light.  The  different  varieties  of  German  silver  contain  different 
proportions  of  copper,  nickel,  and  zinc.  The  per  cent  of  copper  is  50  to 
60,  of  nickel  20  to  25,  and  of  zinc  about  20.  In  color  and  luster  it  re- 
sembles silver,  for  which  it  is  often  substituted.  Its  power  to  conduct 
electricity  is  only  slightly  affected  by  changes  of  temperature,  hence 
it  is  often  used  in  resistance  coils.  Chinese  Pakfong  (or  paktong)  is  a 
variety  of  German  silver.  The  nickel  coins  of  Germany  and  the 
United  States  contain  75  per  cent  copper  and  25  per  cent  nickel.  Cop- 
per is  also  a  constituent  of  many  other  coins.  Britannia  metal  and 
white  metal,  in  which  copper  is  a  minor  constituent,  are  described 
under  Alloys  of  Tin. 

Compounds  of  Copper. — Copper  forms  two  series  of  com- 
pounds, the  cuprous  and  the  cupric.  Thus,  there  are 
cuprous  oxide  (Cu2O)  and  cupric  oxide  (CuO),  cuprous 
chloride  (CuCl)  and  cupric  chloride  (CuCl2).  The  cuprous 
compounds  contain  a  larger  proportion  of  copper  than  the 
cupric  compounds.  Not  every  member  of  each  series  is 
important,  or  even  well  known.  Other  metals  —  mercury 
and  iron  —  form  similar  series.  The  most  important  com- 
pounds are  the  oxides  and  copper  sulphate.  Copper  com- 
pounds are  poisonous.  Cooking  utensils  made  of  copper 
should  be  used  with  care.  Vegetables,  acid  fruits,  and 
preserves,  if  boiled  in  them,  should  be  removed  as  soon  as 
cooked.  The  vessels  themselves  should  be  kept  bright,  to 
prevent  the  formation  of  soluble  copper  salts,  which  might 
contaminate  the  contents. 

Cuprous  Oxide,  Cu2O,  occurs  native  as  cuprite  or  ruby  ore.  It  may 
be  obtained  as  reddish  powder  by  heating  a  mixture  of  solutions  of  cop- 
per sulphate,  Rochelle  salt,  sodium  hydroxide,  and  grape  sugar.  This 
oxide  colors  glass  ruby  red.  It  is  a  beautiful  mineral  and  a  valuable 
ore- 


Copper  —  Silver  —  Gold.  307 

Cupric  Oxide,  CuO,  is  a  black  solid  formed  by  heating 
copper  nitrate.  It  is  reduced  to  metallic  copper  by  hydro- 
gen or  by  carbon,  thus  — 

CuO        +       H2       =     Cu     +   H2O 
Cupric  Oxide         Hydrogen        Copper          Water 

Hence  it  may  be  used  to  determine  the  gravimetric  com- 
position of  water. 

Copper  Sulphate,  CuSO4,  is  the  most  useful  compound 
of  copper.  Like  many  of  the  cupric  compounds  it  is  a  blue, 
crystallized  solid,  and  is  often  called  "  blue  vitriol "  or 
"  blue  stone."  The  crystallized  salt  (CuSO4  .  5  H2O)  loses 
water  in  the  air ;  heated  to  240°  C,  all  the  water  escapes, 
leaving  a  whitish  powder.  This  anhydrous  copper  sul- 
phate absorbs  water  from  alcohol  and  similar  liquids^^nd 
when  added  to  water  it  again  becomes  blue.  Copper -sul- 
phate is  used  in  electric  batteries,  in  making  other  copper 
salts,  in  calico  printing,  dyeing,  copper  plating,  in  preserv- 
ing timber,  and  whenever  a  soluble  copper  compound  is 
needed.  It  is  poisonous  and  is  one  ingredient  of  certain 
mixtures  which  are  sprayed  upon  trees  to  kill  insects. 

Copper  sulphate  may  be  prepared  by  treating  copper 
with  sulphuric  acid.  This  method  is  used  on  a  large  scale, 
but  much  of  the  copper  sulphate  of  commerce  is  a  by- 
product obtained  in  refining  gold  and  silver  with  sulphuric 
acid  (see  below). 

Copper  Nitrate,  Cu(NO3)2,  is  a  blue,  crystallized  solid,  formed  by  the 
interaction  of  copper  and  dilute  nitric  acid.  It  is  a  cupric  salt.  It  is 
very  soluble  in  water,  and  is  readily  decomposed  by  heat  into  cupric 
oxide  and  oxides  of  nitrogen. 

Cuprous  Sulphide,  Cu2S,  is  the  bluish  black  mineral  chalcocite.  Cu- 
pric sulphide,  CuS,  is  the  black  precipitate  formed  by  passing  hydrogen 
sulphide  gas  into  a  solution  of  a  cupric  salt. 


308  Descriptive  Chemistry. 

Malachite  is  a  bright  green  mineral  and  is  often  used  as  an  orna- 
mental stone.  Azurite  is  a  magnificent  blue,  crystallized  mineral. 
Both  are  carbonates  and  valuable  ores  of  copper. 

SILVER. 

Silver  is  one  of  the  precious  metals.  From  the  remotest 
ages  it  has  been  used  for  ornaments,  household  vessels, 
and  money. 

The  Latin  name  of  silver  is  argentine,  from  which  the  symbol  Ag  is 
derived.  The  alchemists  called  it  luna,  on  account  of  its  silvery  or 
"  moonlike"  appearance. 

Occurrence  of  Silver.  —  Native  silver  is  found  in  Ari- 
zona, Mexico,  Norway;  also  in  South  America  and  Aus- 
tralia. The  chief  ores  are  the  sulphides.  The  simple 
sulphide  (silver  glance,  argentite,  Ag2S)  is  the  richest  ore 
and  is  found  in  many  localities  in  the  United  States.  Sil- 
ver sulphide  is  often  combined  with  sulphides  of  lead, 
copper,  antimony,  or  arsenic.  These  complex  sulphides 
are  found  in  Mexico,  Peru,  Bolivia,  Chili,  and  in  Idaho. 
Small  quantities  of  native  silver  chloride  (horn  silver, 
AgCl)  are  also  found ;  it  resembles  wax  or  horn,  and  melts 
in  a  candle  flame.  Sea  water  contains  traces  of  silver,  the 
total  quantity  in  the  ocean  being  estimated  to  be  about  two 
million  tons.  Alloys  of  silver  with  gold,  mercury,  and 
copper  are  found ;  average  California  gold  contains  about 
12  per  cent  silver.  Many  ores  contain  silver,  especially 
those  of  lead;  and  this  argentiferous  (or  silver-bearing) 
lead  is  one  of  the  chief  sources  of  silver. 

The  world's  supply  of  silver  comes  mainly  from  the  United  States. 
Mexico,  Germany,  Australia,  and  Bolivia.  The  United  States  produced 
over  sixty-four  million  ounces  in  1902.  This  was  about  one  third  of  the 
world's  supply,  and  also  the  average  annual  output  for  the  last  few 
years.  Of  this  vast  quantity,  about  90  per  cent  was  furnished  by  Colo- 
rado, Montana,  Utah,  Idaho,  California,  and  Nevada  (Fig.  58). 


Copper  —  Silver  —  Gold. 


309 


Metallurgy  of  Silver.  —  Silver  is  extracted  from  its 
ores  by  two  principal  processes,  (i)  In  the  amalgama- 
tion process  the  powdered  ore  is  first  changed  into  silver 


FIG.  58.  —  Distribution  of  silver  and  gold  in  the  United  States. 

chloride  by  roasting  (or  simply  mixing)  it  with  sodium 
chloride.  The  mass  is  then  reduced  to  silver  by  agitation 
with  water  and  iron  (or  an  iron  compound) ;  the  simplest 
equation  for  this  reaction  is  — 


2  AgCl       + 
Silver  Chloride 


Fe     = 

Iron 


2Ag 
Silver 


FeCl2 

Iron  Chloride 


The  silver  is  removed  by  adding  mercury,  which  forms  an 
amalgam  (an  alloy)  with  the  silver,  but  not  with  the  other 
substances.  When  the  amalgam  is  heated,  the  mercury 
distils  off,  and  the  silver  —  with  some  gold  —  remains  be- 
hind. (2)  Silver  is  extracted  from  lead  ores  by  the  Parkes 
process.  After  the  sulphur,  arsenic,  and  other  impurities 
have  been  removed  from  the  lead  ores,  the  final  product  is 
a  mixture  of  lead,  silver,  and  gold.  This  is  melted  and 


310  Descriptive  Chemistry. 

thoroughly  mixed  with  zinc.  As  the  mixture  cools,  an 
alloy  of  silver,  gold,  zinc,  and  a  little  lead  rises  to  the  top, 
solidifies,  and  is  removed.  The  remaining  lead  mixture  is 
treated  again  with  zinc.  The  alloy  of  silver,  gold,  zinc, 
and  lead  is  heated  to  volatilize  the  zinc  and  to  oxidize  (or 
melt  away)  the  lead.  The  mixture  of  silver  and  gold  is 
heated  with  sulphuric  acid ;  the  gold  is  not  acted  upon, 
but  the  silver  forms  silver  sulphate,  which  is  reduced  by 
copper  to  metallic  silver  (Fig.  59). 


199-7    7 


FIG.  59.  —  Bar  or  "  brick  "  of  silver  showing  the  stamp  of  the  United  States  Assay 
Office  as  a  guarantee  of  its  purity. 


—    — o  —        — r 
Office  as  a  guarantee  of  its  purity 

Lead  ores  containing  considerable  silver  are  sometimes  subjected  to 
CUpellation  to  extract  the  silver.  The  ore  or  alloy  is  heated  in  a  fur- 
nace having  a  shallow  hearth  made  of  porous,  infusible  bone  ash.  The 
lead  is  changed  into  an  oxide  (litharge),  which  melts,  and  is  partly 
driven  off  by  the  air  blast  into  pots  and  partly  absorbed  by  the  porous 
cupel.  The  silver  is  protected  from  the  oxidizing  power  of  the  air  by 
the  melted  litharge,  but  toward  the  end  of  the  operation  the  thin  film  of 
litharge  bursts,  and  the  metallic  silver  appears  as  a  bright  disk  if  the 
operation  is  conducted  in  a  furnace,  and  as  a  globule  or  button  if  the 
extraction  is  performed  in  a  small  assay  cupel.  The  process  is  then 
stopped  and  the  silver  removed. 

Properties  of  Silver.  —  Silver  is  a  lustrous,  white  metal, 
which  takes  a  brilliant  polish.  It  is  harder  than  gold,  but 
softer  than  copper.  Like  copper,  it  is  ductile  and  malle- 
able, and  may  be  easily  made  into  various  shapes.  Its 
specific  gravity  is  about  10.5,  being  heavier  than  copper, 


Copper  —  Silver  —  Gold.  311 

but  lighter  than  lead.  It  melts  at  about  954°  C,  and  fuses 
readily  on  charcoal  in  the  blowpipe  flame ;  it  vaporizes  in 
the  oxyhydrogen  flame  and  in  the  electric  furnace.  Molten 
silver  absorbs  about  twenty  times  its  volume  of  oxygen, 
which  is  expelled  violently  when  the  silver  solidifies.  Pure 
silver  conducts  heat  and  electricity  better  than  any  other 
metal,  but  it  is  too  costly  for  such  uses.  It  does  not  tarnish 
in  air,  unless  sulphur  compounds  are  present,  and  then  the 
familiar  black  film  of  silver  sulphide  is  produced.  This 
blackening  is  especially  noticed  on  silver  spoons  which 
have  been  put  into  eggs  or  mustard,  and  on  silver  coins 
which  have  been  carried  in  the  pocket,  the  sulphur  in  the 
latter  case  coming  from  sulphur  compounds  in  the  perspira- 
tion ;  the  tarnishing  of  household  silver  is  due  to  sulphur 
compounds  in  illuminating  gas  or  gas  from  burning  coal. 
So-called  "  oxidized  "  silver  is  not  oxidized,  but  coated  with 
silver  sulphide.  Silver  is  only  very  slightly  acted  upon  by 
hydrochloric  acid,  and  not  at  all  by  molten  caustic  potash, 
soda,  or  potassium  nitrate.  Nitric  acid  and  hot  concen- 
trated sulphuric  acid  change  it  into  the  nitrate  and  sulphate, 
respectively,  as  in  the  case  of  copper. 

Alloys  of  Silver.  —  Pure  silver  is  too  soft  for  constant 
use,  and  is  usually  hardened  by  adding  a  small  amount  of 
copper.  These  alloys  are  used  as  coins  and  for  jewelry. 
The  silver  coins  of  the  United  States  and  France  contain 
900  parts  of  silver  to  100  of  copper,  and  are  called  900 
fine.  British  silver  coins  are  925  fine ;  this  quality  is  called 
"sterling  silver,"  and  from  it  much  ornamental  and  useful 
silverware  is  made. 

Silver  Plating.  —  Metals  cheaper  than  silver  may  be 
coated  or  plated  with  pure  silver  precisely  as  in  the  case  of 
copper.  Plated  silverware  has  the  appearance  of  solid  or 
pure  silver.  The  object  to  be  plated  is  carefully  cleaned, 


312  Descriptive  Chemistry. 

and  made  the  cathode  in  a  bath  or  solution  of  potassium 
silver  cyanide.  The  anode  is  a  plate  of  pure  silver  (Fig. 

60).  The  deposit  of  silver  is 
dull,  but  may  be  brightened 
by  rubbing  with  or  without 
chalk. 

Compounds    of    Silver.  - 
FIG.  60.  —  Apparatus  for  silver  plat-    The  most  important  compound 

ing.     A,  A,  A,  are  silver  anodes,  and      .  .,  . ,  /  .      AT ^  x 

thf  spoons  are  cathodes.  «     SllvCI     nitrate     (AgNO3). 

It  is  a  white  crystalline  solid, 

made  by  dissolving  silver  in  nitric  acid.  Exposed  to  the 
light,  it  turns  dark  if  in  contact  with  organic  matter.  It 
discolors  the  skin ;  if  applied  long  enough,  it  disintegrates 
the  flesh,  and  is  often  used  by  physicians  for  this  purpose. 
Its  caustic  action  and  the  silvery  color  of  the  metal  from 
which  it  is  made  long  ago  led  to  its  name,  lunar  caustic. 
Besides  its  extensive  use  in  photography  and  silver  plating, 
silver  nitrate  is  the  essential  constituent  of  indelible  ink. 
Silver  chloride  (AgCl)  is  made  by  adding  hydrochloric 
acid  or  the  solution  of  any  chloride  to  a  solution  of  a  silver 
compound.  Thus  formed,  it  is  a  white,  curdy  solid,  which 
turns  violet  in  the  light,  and  finally  black.  This  action  of 
light  is  more  intense  if  organic  ntatter  is  present.  It 
dissolves  in  ammonium  hydroxide,  forming  a  complex  com- 
pound of  the  two  substances.  The  formation  and  proper- 
ties of  silver  chloride  constitute  the  test  for  silver.  Silver 
bromide  (AgBr)  and  silver  iodide  (Agl)  are  analogous  to 
silver  chloride  in  their  properties  and  methods  of  forma- 
tion. They  are  used  in  photography. 

Photography  is  based  on  the  fact  that  silver  salts,  espe- 
cially the  bromide  and  iodide,  change  color  when  mixed 
with  organic  matter  and  exposed  to  the  light.  The  photo- 
graph is  taken  on  a  glass  plate,  coated  on  one  side  with  a 


Copper  —  Silver  —  Gold.  313 

thin  layer  of  gelatine,  containing  the  silver  salts.  Some- 
times a  sheet  of  sensitized  gelatine,  called  a  film,  is  used. 
The  plate  or  film  is  placed  in  the 'camera  and  exposed. 
The  light,  which  comes  from  the  object  being  photographed, 
changes  the  silver  salts  in  proportion  to  its  brilliancy.  The 
plate,  however,  shows  no  change  until  it  has  been  devel- 
oped. This  process  consists  in  treating  the  plate  with  a 
reducing  agent,  e.g.  ferrous  sulphate,  pyrogallic  acid,  or 
special  mixtures.  As  the  developer  acts  upon  the  plate, 
the  image  appears.  This  is  really  a  deposit  of  finely 
divided  silver.  Where  the  intense  light  fell  upon  the 
plate,  the  deposit  is  heavier  than  where  little  or  no  light 
fell.  Hence,  dark  parts  of  the  object  appear  light  on  the 
plate,  and  light  parts  dark;  and  since  the  image  is  the 
reverse  of  the  object,  the  plate  is  called  a  negative.  When 
the  plate  has  been  properly  developed,  it  still  contains  sil- 
ver salts  not  altered  by  the  light ;  and  if  they  were  left  on  the 
plate,  the  image  would  be  clouded,  and  finally  obliterated 
by  the  light.  The  image  is,  therefore,  fixed  by  wash- 
ing off  the  silver  salts  with  a  solution  of  sodium  thiosul- 
phate  (or  "hyposulphite").  A  print  is  made  by  laying 
sensitized  paper  upon  the  negative  and  exposing  them  to 
the  sunlight,  so  that  the  light  will  pass  through  the  nega- 
tive. The  negative  obstructs  the  light  in  proportion  to  the 
thickness  of  the  silver  deposit,  so  the  photograph  has  the 
same  shading  as  the  object.  Most  prints,  like  the  plates, 
must  be  fixed.  Sometimes  the  color  is  improved  by  toning, 
i.e.  by  placing  the  print  in  a  solution  of  gold  or  of  platinum. 

GOLD. 

Gold  is  the  most  precious  of  the  metals,  and  has 
been  used  from  the  earliest  times  for  adornment  and  as 
money. 


314  Descriptive  Chemistry. 

The  Latin  name  of  gold,  aurum,  gives  the  symbol  Au.  For  many 
centuries  the  alchemists  tried  to  produce  gold  from  base  or  cheaper 
metals.  They  were  unsuccessful  in  their  search  for  the  Philosopher's 
Stone,  which  they  believed  had  power  to  affect  this  transformation. 

Occurrence  of  Gold.  —  Gold  is  widely  distributed,  but 
not  abundantly  in  many  places.  Unlike  copper  and  silver, 
its  compounds  are  few  and  rare;  the  only  important  ones 
are  the  tellurides  (compounds  of  tellurium)  found  in  Colo- 
rado. It  is  never  found  pure,  being  alloyed  with  silver 
and  occasionally  with  copper  or  iron.  It  is  disseminated 
in  fine,  almost  invisible,  particles  among  ores  of  other 
metals,  though  not  so  abundantly  as  silver.  Much  gold  is 
found  in  veins  of  quartz,  and  in  the  sand  and  gravel  formed 
from  gold-bearing  rocks.  Gold  occurs  usually  as  dust, 
scales,  or  grains,  but  occasionally  shapeless  masses  called 
"  nuggets  "  are  found,  varying  in  weight  from  a  few  grams 
to  many  kilograms.  The  largest  nugget  ever  known 
weighed  over  84  kg.  (184  Ibs.). 

The  chief  gold-producing  countries  are  the  United  States,  Australia, 
South  Africa,  and  Russia.  In  1902  the  United  States  produced  nearly 
four  million  ounces,  which  came  largely  from  Colorado,  California,  and 
other  Western  states,  and  Alaska.  Gold  in  working  quantities  is  found 
in  about  twenty  states  of  the  Union  (Fig.  58).  The  total  value  of  the 
gold  produced  in  the  world  in  1902  was  about  $306,000,000. 

Gold  Mining.  —  Gold  was  first  obtained  by  miners  by 
washing  the  gold-bearing  sand  and  gravel  of  a  stream  in 
large  pans  or  cradles.  This  primitive  method  was  soon  re- 
placed by  placer  mining  and  hydraulic  mining.  Streams 
of  water,  directed  against  the  earth  containing  the  gold, 
wash  away  the  lighter  materials,  but  leave  the  heavy  gold 
behind  in  the  form  of  scales  or  "  gold  dust."  From  this 
mixture  gold  and  silver  are  extracted  by  mixing  with  mer- 
cury, or  by  passing  the  moistened  mass  over  copper  plates 


Copper  —  Silver  —  Gold.  315 

coated  with  mercury.  The  amalgam  is  then  heated,  as  in 
the  metallurgy  of  silver,  to  remove  the  mercury ;  the  resi- 
due of  gold  and  silver  is  purified  as  described  below.  In 
vein  mining  the  gold-bearing  rock  —  usually  quartz  —  is 
crushed  and  then  washed,  and  the  gold  removed  by  mer- 
cury, as  in  placer  mining  (see  Chapter  XX).  Low  grade 
ores  and  those  containing  certain  metals  cannot  be  profita- 
bly treated  with  mercury.  In  the  chlorination  process  the 
crushed  ore  is  roasted  and  then  revolved  in  barrels  contain- 
ing bleaching  powder  and  sulphuric  acid;  this  operation 
forms  a  soluble  gold  chloride  ( AuCl3),  from  which  the  gold 
is  precipitated  as  a  fine  powder  by  hydrogen  sulphide  (or 
other  reducing  agents).  In  the  cyanide  "process  the 
crushed  ore,  or  the  slime  from  a  previous  extraction,  is 
mixed  with  a  weak  solution  of  potassium  cyanide  and 
exposed  to  the  air ;  this  operation  changes  the  gold  into  a 
soluble  cyanide  (KAu(CN)2).  The  gold  is  separated  from 
this  solution  by  electrolysis  or  by  treatment  with  zinc. 

Purification  of  Gold.  —  Gold  obtained  by  the  above 
methods  is  impure,  silver  being  the  chief  impurity.  These 
metals  are  parted  by  a  chemical  process  or  separated  by 
electrolysis.  By  the  old  parting  process  known  as  quar- 
tation  an  alloy  of  gold  and  silver,  in  which  the  gold  is 
about  one  fourth  of  the  whole,  is  treated  with  nitric  acid ; 
this  operation  changes  the  silver  into  the  nitrate  from 
which  the  pure  gold  may  be  readily  removed.  The  metals 
may  be  parted  by  the  cheaper  method  described  under 
silver,  viz.  by  boiling  with  concentrated  sulphuric  acid. 
By  this  treatment  the  gold,  which  is  about  one  sixth  of  the 
alloy,  is  left  as  a  brownish,  porous  mass.  It  is  washed, 
dried,  and  fused  with  charcoal  and  sodium  carbonate.  In 
the  electrolytic  method  of  separation,  the  anode  is  an 
alloy  of  gold  and  silver,  the  cathode  is  silver,  and  the  elec- 


3 1 6  Descriptive  Chemistry. 

trolyte  is  nitric  acid.  When  the  current  passes,  part  of  the 
silver  of  the  anode  goes  into  solution  as  the  nitrate,  while 
part  is  deposited  at  the  cathode ;  the  gold  remains  at  the 
anode  as  a  fine  powder  and  is  caught  in  a  cloth  bag  which 
incloses  the  whole  anode.  Gold  is  now  purified  at  the 
United  States  Mint  by  electrolysis.  The  electrolyte  is  a 
hydrochloric  acid  solution  of  gold  chloride,  but  otherwise 
the  process  is  the  same  as  described  above.  It  is  more 
economical  than  parting  by  nitric  acid. 

The  purity  of  gold  is  expressed  in  carats.  Pure  gold  is 
24  carats  fine ;  an  alloy  containing  22  parts  of  gold  and  2 
parts  copper  is  22  carat  gold,  while  one  containing  equal 
parts  gold  and  other  metals  is  12  carat  gold  (see  foot-note, 
page  183). 

Properties  of  Gold.  —  Gold  is  a  yellow  metal.  It  is 
about  as  soft  as  lead,  and  is  the  most  ductile  and  malleable 
of  all  metals.  The  leaf  into  which  it  may  be  beaten  is 
very  thin  and  is  green  by  transmitted  light.  Air,  oxygen, 
and  most  acids  do  not  attack  it ;  but  it  is  changed  into  a 
gold  chloride  (AuCl3)  by  aqua  regia  (see  Aqua  Regia). 
Gold  is  one  of  the  heaviest  metals,  its  specific  gravity 
being  about  19. 

Uses  of  Gold.  —  Pure  gold  is  too  soft  for  most  practical 
purposes,  and  is,  therefore,  usually  hardened  with  copper 
or  silver.  The  gold-copper  alloy  has  a  reddish  color  and 
is  often  called  "  red  gold  "  ;  the  gold-silver  alloy  is  paler 
than  pure  gold  and  is  sometimes  called  "white  gold." 
Gold  coins  contain  gold  and  copper.  The  United  States 
standard  gold  coins  contain  9  parts  gold  and  I  part  cop- 
per, while  in  England  the  legal  standard  is  n  of  gold  to 
i  of  copper.  Gold  leaf  of  various  grades  is  used  to  orna- 
ment books,  signs,  and  many  objects.  Jewelers  use  gold 
for  many  purposes;  such  gold  varies  from  1 2  to  22  carats 


Copper  —  Silver  —  Gold.  317 

in  purity.  On  account  of  its  malleability,  feeble  chemical 
action,  and  beauty,  gold  is  used  by  dentists  for  filling 
teeth. 

Compounds  of  Gold  are  readily  decomposed  by  metals,  weak  reducing 
agents  (e.g.  ferrous  sulphate  or  hydrogen  sulphide),  fine  solids  like  char- 
coal, and  by  electrolysis.  When  gold  is  dissolved  in  aqua  regia  and 
the  acid  removed  by  evaporation,  the  resulting  gold  chloride  (AuCl3) 
gives  with  stannous  chloride  solution  a  beautiful  purple  precipitate  ;  the 
latter  is  called  "purple  of  Cassius,"  and  is  probably  finely  divided  gold. 
Its  formation  is  the  test  for  gold.  The  process  of  gold  plating  is  the 
same  as  silver  plating,  only  the  solution  is  one  of  potassium  gold  cyan- 
ide (Au(CN)3 .  KCN)  and  the  anode  is  gold.  Much  cheap  jewelry  is 
gold  plated. 

EXERCISES. 

1.  What  is  the  symbol  of  (a}  copper,  (b)  silver,  (c}  gold?     State 
the  derivation  of  each  symbol. 

2.  Where  is  copper  found  abundantly?     State  in   what  form   it 
occurs  in  each  locality.     Discuss  its  production. 

3.  Describe  briefly  the  metallurgy  of  (a)  native  copper,  (£)  oxides 
and  carbonates,  (c)  copper-iron  sulphides. 

4.  Describe  the  purification  of  copper  by  electrolysis. 

5.  State    (#)    the    physical    properties   of   copper,   and    (£)    the 
chemical  properties. 

6.  Describe  several  tests  for  copper. 

7.  Discuss  the  uses  of  copper. 

8.  Name  ten  alloys  of  copper.     Describe  five  important  alloys. 

9.  What  is  an  electrotype?     How  is  it  made?     (See  Chapter  X.) 

10.  State  the  general  properties  of  copper  compounds. 

11.  Describe  the  oxides  of  copper. 

12.  Describe  the  manufacture,  and  state  the  properties  and  uses  of 
copper  sulphate. 

13.  What  are  the  properties  of  (a)  copper  nitrate,  (b}  malachite, 
and  (<:)  azurite? 

14.  What  is  the  formula  of  (#)  copper  sulphate,  ($)  copper  nitrate, 
(c)  cupric  oxide,  and  (*/)  cuprous  oxide?  / 

15.  Discuss  (#)  the  occurrence,  and  (<£)  the  production  of  silver. 

1 6.  Describe  the  extraction  of  silver  by  (a)  the  amalgamation  pro- 
cess, and  (£)  the  Parkes  process. 


318  Descriptive  Chemistry. 

17.  State  (a)  the  physical  properties  of  silver,  and  (£)  the  chemical 
properties. 

1 8.  Discuss  (#)  silver  alloys,  and  (b)  silver  plating. 

19.  State  the  properties  and  uses  of  silver  nitrate. 

20.  State  the  properties  of  silver  chloride.     What  is  the  test  for 
silver? 

21.  Describe  briefly  the  essential  operations  in  photography.    What 
general  chemical  changes  does  it  utilize? 

22.  What  is  (a}   blue  vitriol.  (<£)  argentiferous  lead,  (c)   oxidized 
silver,  (d)   sterling  silver,  (e)  coin  silver,  (/)   lunar  caustic,  and  (g) 
"hypo"? 

23.  What  is  the  formula  of  (a)  silver  nitrate,  and  (£)  silver  chloride? 

24.  Discuss  (a)  the  occurrence  of  gold,  and  ($)  its  production. 

25.  Describe  the  different  methods  of  (a)  mining,  and  (£)  extracting 
gold. 

26.  Describe  the  purification  of  gold  by  (a)  parting,  and  (£)  elec- 
trolysis. 

27.  What  is  1 8  carat  gold? 

28.  State  (a}  the  properties,  and  (/£)  the  uses  of  gold. 

29.  Discuss  (#)  compounds  of  gold,  and  (b}  gold  plating. 

30.  What  is  the  test  for  gold? 

31.  What  is   (a)  gold  dust,  (<£)  aqua  regta,  (V)  a  nugget,  (d)  gold 
leaf? 

PROBLEMS. 

1.  How  much  cupric  oxide  is  formed  by  heating  1467  gm.  of  copper 
in  air?     (Assume  Cu  +  O  =  CuO.) 

2.  Calculate  the  per  cent  of  copper  in  (a)  malachite   (CuCO» . 
Cu(OH)9),    (£)    azurite    (2  CuCO3 .  Cu(OH)9),    (c)   copper  sulphate 
(CuS04). 

3.  If  480  gm.  of  silver  interact  with  nitric  acid,  how  much  silver 
nitrate  is  formed? 

4.  Calculate  the  per  cent  of  silver  in  (a}  silver  chloride  (AgCi), 
(£)  silver  sulphide  (AgS2),  (c)  silver  nitrate  (AgNO3). 


CHAPTER    XXIII. 
CALCIUM,   STRONTIUM,  AND  BARIUM. 

THESE  elements  form  a  natural  group  called  the  alkaline 
earth  metals.  The  metals  themselves  are  rare,  but  their 
compounds,  especially  those  of  calcium,  are  numerous  and 
useful.  This  group  resembles  the  alkali  group. 

CALCIUM. 

Occurrence  of  Calcium. — Calcium  is  never  found  free. 
Combined  calcium  makes  up  about  3.5  per  cent  of  the 
earth's  crust.  The  most  abundant  compound  is  calcium 
carbonate  (CaCO3).  This  has  many  familiar  forms,  e.g. 
limestone,  chalk,  marble,  coral,  and  shells.  Many  rocks 
are  complex  silicates  of  calcium  and  other  metals.  The 
extensive  deposits  of  calcium  phosphate,  calcium  borate, 
and  calcium  fluoride  have  been  mentioned.  Calcium  sul- 
phate (CaSO4)  occurs  abundantly  in  the  form  of  gypsum, 
alabaster,  and  selenite.  Calcium  compounds  are  essential 
to  the  life  of  plants  and  animals,  being  found  in  the  leaves 
of  plants,  and  in  the  bones,  teeth,  and  shells  of  animals. 
Many  rivers  and  springs  contain  calcium  salts,  especially 
the  acid  carbonate  and  sulphate. 

Preparation  and  Properties.  —  Metallic  calcium  was  obtained  by 
electrolysis  in  1808  by  Davy,  but  our  knowledge  of  the  pure  metal  is 
due  to  Moissan.  In  1898  he  prepared  it  from  the  iodide  by  electrolysis 
and  by  fusion  with  sodium.  The  equation  for  the  latter  process  is  — 

CaI2  +    2  Na     =     Ca       +  2  Nal 

Calcium  Iodide      Sodium      Calcium      Sodium  Iodide 


320 


Descriptive  Chemistry. 


Calcium  is  a  silver-white  metal,  soft  enough  to  be  cut  with  a  knife, 
though  harder  than  potassium.  It  may  be  crystallized  from  melted 
sodium.  It  readily  decomposes  water  at  the  ordinary  temperature,  and 
combines  directly  with  most  of  the  other  elements. 

Calcium  Carbonate,  CaCO3. — The  most  abundant  form 
of  this  compound  is  limestone.  Vast  deposits  are  found 
in  many  places,  exhibiting  a  variety  of  textures  and  colors. 
In  the  United  States  much  limestone  is  found  in  Iowa, 
Illinois,  and  Wisconsin.  All  kinds  are  compact  and  usually 
soft,  though  some  are  hard  enough  for  use  as  building 
stone;  some  are  coarse,  and  often  consist  of  grains,  crys- 
tals, or  small  shells.  Pure  limestone  is  white  or  gray,  but 

impurities,  especially  organic 
matter  and  iron  compounds,  pro- 
duce blue,  yellow,  reddish,  and 
black  varieties.  Hard,  crystal- 
line limestone  which  takes  a  good 
polish  is  called  marble.  This 
form,  which  has  a  wide  range  of 
color,  is  used  as  a  building  and 
an  ornamental  stone.  Calcite 


FIG.  61.  —  Calcite  crystals. 


is  crystallized  calcium  carbonate.     It  is  almost  as  abundant 

as      quartz,     though 

softer ;      its      varied 

color  and  crystal  form 

combine   to  make  it 

attractive   (Fig.   61). 

A    very    transparent 

variety      of      calcite 

called   Iceland   spar 


FlG.  62.  —  Crystallized  Iceland  spar  showing 
double  refraction. 


has  the  remarkable 
property  of  double 
refraction,  i.e.  of  making  objects  appear  double  (Fig.  62). 


Calcium,  Strontium,  and  Barium.  321 

Calcium  carbonate  is  not  soluble  in  water,  unless  carbon 
dioxide  is  present  (see  Carbon  Dioxide).  As  water  con- 
taining carbon  dioxide  works  its  way  underground  in 
limestone  regions  the  limestone  is  dissolved  and  caves 


FIG.  63.  —  Stalactites  and  stalagmites  in  Luray  Cavern. 
From  a  photograph  copyrighted  by  C.  H.  James. 

are  often  formed  or  enlarged.  When  the  water  enters  a 
cave  and  drips  from  the  top,  the  water  evaporates,  or 
the  gas  escapes,  or  both,  and  the  calcium  carbonate  is 
redeposited,  often  forming  stalactites  and  stalagmites 
(Fig.  63).  The  stalactites  hang  from  the  roof  like  icicles, 


322  Descriptive  Chemistry. 

while  the  stalagmites  grow  up  from  the  floor,  as  the 
deposit  slowly  accumulates  from  the  solution  which 
drops  from  the  roof  or  the  tips  of  stalactites.  The 
Mammoth  Cave  in  Kentucky,  the  Marengo  Cave  in  Indi- 
ana, and  the  Luray  Cavern  in  Virginia  are  famous  for 
these  fantastic  formations.  Mexican  onyx  is  a  variety  of 
stalagmite.  Vast  deposits  of  this  beautiful  mineral  are 
found  in  Algeria  and  Mexico.  It  is  translucent  and  deli- 
cately colored,  and  is  used  as  an  ornamental  stone,  espe- 
cially for  altars,  table  tops,  mantels,  and  lamp  standards. 
Beautiful  deposits  of  limestone  are  found  around  many 
mineral  springs.  Travertine  occurs  near  many  springs 
in  Italy.  When  fresh,  it  is  soft  and  porous,  but  it  soon 
hardens  and  becomes  a  durable  building  stone  in  dry  cli- 
mates. The  outer  walls  of  the  Colosseum  and  of  St. 
Peter's  are  travertine.  Limestone  often  contains  shells 
and  fossils,  confirming  our  belief  that  limestone  is  the 
remains  largely  of  the  shells  of  animals.  The  calcium 
carbonate  dissolved  in  the  ocean  is  transformed  by  marine 
organisms  into  shells  and  bony  skeletons.  The  hard  parts 
of  these  animals  accumulate  in  vast  quantities  on  the  ocean 
bottom,  become  compact,  often  hardened  and  crystallized, 
and  are  finally  elevated  into  their  present  position.  On 
the  coast  of  Florida,  coquina  or  shell  rock  is  found.  It  is 
a  mass  of  fragments  of  shells  cemented  by  calcium  carbon- 
ate, and  in  time  will  become  compact  limestone.  Chalk 
is  the  remains  of  shells  of  minute  animals.  When  exam- 
ined under  a  microscope,  a  good  specimen  is  seen  to  con- 
sist almost  entirely  of  tiny  shells.  The  ocean  contains 
myriads  of  minute  animals,  and  when  they  die,  their  shells, 
which  are  calcium  carbonate,  sink  to  the  bottom.  As  a 
result,  the  ocean  bottom  is  partly  covered  with  a  gray 
mud,  called  globigerina  ooze.  Under  the  microscope  this 


Calcium,  Strontium,  and  Barium. 


3*3 


ooze  looks  like  Figure  64,  and  when  dried  and  compressed 
it  can  hardly  be  distinguished  from  chalk.  Hence  it  is  be- 
lieved that  the  immense  beds  of  chalk  found  in  England 
and  other  places  were  formed  from  this  ooze.  Some  vari- 
eties of  chalk  under  the  microscope  resemble  the  ooze 


FIG.  64. —  Ooze  from  the  ocean 
bottom,  showing  globigerina  shells 
(magnified). 


FIG.  65.  —  Chalk  from  Iowa,  showing 
globigerina  shells  (magnified). 


(Fig.  65).  Blackboard  crayon  is  a  mixture  of  chalk  and 
clay.  Whiting  is  a  variety  of  impure  chalk ;  putty  is  a 
mixture  of  whiting  and  oil.  Coral  is  calcium  carbonate. 
The  vast  accumulations  in  the  sea  are  the  skeletons  of  the 
coral  animals. 

The  properties  of  calcium  carbonate,  discussed  in  Chap- 
ter XIV,  may  be  profitably  reviewed  at  this  point. 

Besides  being  burned  into  lime,  immense  quantities  of 
limestone  are  consumed  in  manufacturing  iron  and  steel, 
the  United  States  alone  using  annually  over  seven  million 
tons  in  this  industry. 


324  Descriptive  Chemistry. 

Calcium  Oxide,  CaO,  is  the  chemical  name  of  lime.  It 
is  a  hard,  white  solid.  Pure  lime  is  almost  infusible,  and 
when  heated  in  the  oxyhydrogen  flame,  it  gives  an  in- 
tensely bright  light,  sometimes  called  the  "lime  light"  (see 
Hydrogen).  In  the  electric  furnace  it  melts  and  volatil- 
izes, if  the  heating  is  prolonged.  Lime  containing  impuri- 
ties, like  sand,  clay,  and  iron  compounds,  melts  quite  readily 
into  a  glass  or  slag.  Exposed  to  the  air,  lime  becomes 
"  air  slaked,"  i.e.  it  slowly  absorbs  water  and  carbon  diox- 
ide, swells,  and  soon  crumbles  to  a  powder,  which  is  a 
mixture  of  calcium  hydroxide  and  calcium  carbonate. 
Lime  and  water  combine  violently  and  liberate  consider- 
able heat,  as  is  often  seen  when  mortar  is  being  prepared. 
This  operation  is  called  "  slaking,"  and  the  product  is 
"slaked  lime."  The  equation  for  the  chemical  change 
is  — 

CaO       +     H2O    =        Ca(OH)2 
Calcium  Oxide         Water        Calcium  Hydroxide 

Fresh  lime  attacks  organic  matter,  and  is  therefore  often 
called  "  caustic  lime  "  or  quicklime.  It  combines  with 
water  to  form  calcium  hydroxide  and  with  acids  to  form 
calcium  salts. 

Lime  is  one  of  the  most  important  substances.  It 
is  used  in  preparing  mortar,  cement,  metals,  in  making 
bleaching  powder,  calcium  carbide,  sodium  hydroxide, 
and  glass,  in  purifying  illuminating  gas  and  sugar,  to 
remove  hair  from  hides  before  the  process  of  tanning,  in 
dyeing  and  bleaching  cotton  cloth,  in  drying  gases,  and  as 
a  disinfectant  and  fertilizer. 

Lime  is  prepared  on  a  large  scale  by  heating  limestone 
in  a  partly  closed  cavity  or  vessel.  The  decomposition 
takes  place  according  to  the  equation — • 


Calcium,  Strontium,  and  Barium. 


325 


FIG.  66.— Limekiln  (ver- 
tical section).  The  fire  is 
built  in  A  under  the  arch  of 
limestone. 


CaCO3       =       CaO        +         CO2 
Calcium  Carbonate     Calcium  Oxide     Carbon  Dioxide 

The  carbon  dioxide  gas  escapes  and  the  lime  is  left  in  the 
kiln. 

Limestone  was  formerly  "burned"  in  a 
cavity  on  a  hillside,  and  in  some  regions  it  is 
so  prepared  to-day.  An  arch  of  limestone  is 
built  across  the  cavity  above  the  fire  pit,  and 
limestone  is  piled  upon  the  arch  until  the 
kiln  is  full  (Fig.  66).  The  fire  is  then  lighted 
and  kept  burning  for  about  three  days.  These 
kilns  have  been  largely  replaced  by  a  modern 
furnace,  constructed  so  that  the  heat  can  be 
regulated,  the  gases  swept  out,  and  the  prod- 
uct removed  without  extinguishing  the  fire. 

Limestone,  containing  more  than  10  per 
cent  of  clay,  forms  hydraulic  lime,  which 

becomes  very  hard  when  wet  or  kept  in  contact  with  water.  Cements 
are  varieties  of  hydraulic  lime.  They  are  made  by  burning  a  mixture 
(natural  or  artificial)  of  limestone,  clay,  and  sand,  and  grinding  the 
product  to  a  very  fine  powder.  Rosendale  and  Portland  are  the 
common  brands.  The  hardening  of  cements  is  not  well  understood. 

Calcium  Hydroxide,  Ca(OH)2,  is  a  white  powder.  It 
is  sparingly  soluble  in  water,  but  more  soluble  in  cold  than 
in  warm  water.  The  solution  has  a  bitter  taste,  an  alkaline 
reaction,  and  is  commonly  called  limewater.  Exposed  to 
the  air,  limewater  becomes  covered  with  a  thin  crust  of  cal- 
cium carbonate,  owing  to  the  absorption  of  carbon  dioxide. 
For  the  same  reason,  limewater  becomes  milky  or  cloudy 
when  carbon  dioxide  is  passed  into  it.  The  formation  of 
calcium  carbonate  in  this  way  is  the  usual  test  for  carbon 
dioxide.  The  equation  for  this  chemical  change  is  — 

Ca(OH)2  +   C02  =   CaC03  +   H2O 
Limewater         Carbon        Calcium 
Dioxide     Carbonate 


326  Descriptive  Chemistry. 

Limewater  is  prepared  by  carefully  adding  lime  to  consid- 
erable water,  allowing  the  mixture  to  stand  until  the  solid 
has  settled,  and  then  removing  the  pure  liquid.  When 
considerable  calcium  hydroxide  is  suspended  in  the  liquid, 
the  mixture  is  called  milk  of  lime.  Ordinary  whitewash 
is  thin  milk  of  lime.  Limewater  is  used  in  the  chemical 
laboratory  and  as  a  medicine. 

Mortar  is  a  thick  paste  formed  by  mixing  lime,  sand,  and  water.  It 
slowly  hardens  or  "  sets/'  owing  to  the  loss  of  water  and  to  the  absorp- 
tion of  carbon  dioxide.  It  hardens  without  much  shrinking,  and  when 
placed  between  bricks  or  stones  holds  them  firmly  in  place.  The  sand 
makes  the  mass  porous  and  thus  facilitates  the  change  of  the  hydroxide 
into  the  carbonate.  The  sand  itself  is  changed  chemically  only  to  a 
slight  extent,  if  at  all.  Hair  is  sometimes  added  to  make  the  mortar 
stick  better,  especially  when  it  is  used  as  plaster  for  walls. 

Calcium  Sulphate,  CaSO4.—  Extensive  deposits  of  the 
different  forms  of  calcium  sulphate  are  found  in  England, 
France,  Nova  Scotia,  and  in  the  United  States,  especially 
in  Michigan,  Kansas,  Iowa,  Virginia,  Tennessee,  and  Ken- 
tucky. It  is  generally  found  in  volcanic  regions,  and -is 
often  associated  with  sulphur  and  limestone,  one  variety 
(anhydrite,  CaSO4)  being  found  with  salt.  Gypsum 
occurs  as  white  masses  or  transparent  crystals,  having  the 
composition  CaSO4  .  2  H2O.  Lustrous,  translucent,  soft 
crystals  are  called  selenite.  Fine  grained,  massive  kinds 
are  known  as  alabaster,  and  the  fibrous  kinds  as  satin 
spar. 

Gypsum  is  widely  used  as  a  fertilizer,  and  in  making 
glass  and  porcelain.  Alabaster,  being  soft  and  beautiful, 
is  carved  into  statues  and  other  ornaments. 

Calcium  sulphate,  when  heated,  loses  its  water  of  crys- 
tallization, becomes  opaque,  and  falls  to  a  powder.  This 
powder,  if  moistened,  swells  and  quickly  "  sets  "  or  solidi- 


LA 

*^ 

Calcium,  Strontium,  and  Barium.  327 

fies  to  a  white,  porous  mass  with  a  smooth  surface.  When 
properly  prepared  this  powder  is  plaster  of  Paris,  which 
derives  its  name  from  the  celebrated  gypsum  beds  near 
Paris.  Plaster  of  Paris  is  used  to  coat  walls,  to  cement 
glass  to  metal,  but  more  largely  to  make  casts  and  repro- 
ductions of  statues  and  small  objects.  Stucco  is  essen- 
tially a  mixture  of  glue  and  plaster  of  Paris. 

To  make  plaster  of  Paris,  lumps  of  gypsum  (CaSO4 .  2  H2O)  are 
heated  to  about  125°  C.  to  expel  part  of  the  water.  The  product 
((CaSO4)2 .  H2O)  is  ground  fine.  The  "  setting  "  is  a  chemical  change. 
The  slightly  soluble  plaster  of  Paris  slowly  combines  with  water  to  form 
a  network  of  very  small  crystals  of  the  less  soluble  hydrated  calcium 
sulphate.  The  equation  is — 

(CaS04)2.H2O   +   3H2O   =   2(CaSO4 .  2  H2O) 
Plaster  of  Paris         Water  Gypsum 

Calcium  Compounds  and  Hardness  of  Water.  —  Calcium 
sulphate  is  slightly  soluble  in  water,  and  calcium  carbon- 
ate, as  we  have  already  seen,  is  changed  into  the  unstable 
acid  carbonate  by  water  containing  carbon  dioxide.  Water 
containing  these  salts  of  calcium  is  called  hard  water. 
They  form  sticky,  insoluble  compounds  with  soap,  and  as 
long  as  water  contains  such  salts,  the  soap  is  useless  as 
a  cleansing  agent.  Heat  decomposes  acid  calcium  carbon- 
ate, and  the  hardness  due  to  calcium  carbonate  is  called 
temporary  hardness,  because  boiling  removes  it.  But 
the  hardness  caused  by  calcium  sulphate  cannot  be  so  re- 
moved, and  is  called  permanent  hardness.  Magnesium 
sulphate,  like  calcium  sulphate,  produces  permanent  hard- 
ness. Soft  water,  such  as  rain  water,  contains  little  or  no 

ilcium  or  magnesium  salts. 

Calcium  Chloride,  CaCl2,  is  a  white  solid.  It  absorbs 
moisture  rapidly,  and  is  used  to  dry  many  gases  and 

<;uids.       The    crystallized   variety    dissolves    readily   in 


328  Descriptive  Chemistry. 

water,  and  the  solution  is  attended  by  a  marked  fall  of 
temperature.  A  mixture  of  crystallized  calcium  chloride 
and  snow  produces  a  temperature  of  —  40°  C.  The  liquid 
left  from  the  interaction  of  calcium  carbonate  and  hydro- 
chloric acid  contains  calcium  chloride,  which  on  concentra- 
tion is  deposited  in  large  crystals.  These  readily  absorb 
water,  but  lose  their  own  water  of  crystallization  when 
heated  above  200°  C.  This  anhydrous  calcium  chloride  is 
porous,  and  is  the  form  usually  used  as  a  drying  agent. 
At  a  high  temperature  it  melts,  and  solidifies  in  cooling 
to  a  hard  mass  known  as  fused  calcium  chloride. 

Calcium  chloride  is  found  in  small  quantities  in  some  of  the  Stass- 
furt  salts.  It  is  obtained  in  large  quantities  as  a  by-product  in  the 
manufacture  of  sodium  carbonate  (by  the  Solvay  process)  and  other 
chemicals. 

Other  Compounds  of  Calcium  have  already  been  discussed  and  may 
be  reviewed  here.  They  are  calcium  fluoride,  calcium  carbide,  the  cal- 
cium phosphates,  and  calcium  hypochlorite.  Calcium  sulphide  (CaS) 
is  formed  by  heating  a  mixture  of  gypsum  and  carbon ;  like  other  sul- 
phides, it  stains  silver  brown. 

Test  for  Calcium.  —  Calcium  compounds,  especially  the  chloride, 
color  the  Bunsen  flame  a  yellowish  red. 

STRONTIUM    AND    BARIUM. 

Strontium,  Sr,  and  Barium,  Ba,  are  uncommon  metallic  elements. 
They  resemble  calcium  closely  in  their  physical  properties  and  chem: 
relations.      The  metals  themselves   never  occur  free,  and  are  har  !! 
more  than  chemical  curiosities.     Their  compounds  are  abundant,  and 
some  are  useful. 

Compounds  of  Strontium.  —  The  important  native  compounds  are 
the  beautifully  crystallized  minerals,  strontianite  (strontium  carbor  . 
SrCO3)  and  celestite  (strontium  sulphate,  SrSO4).  Strontium  oxidf 
(strontia,  SrO),  like  Ihne,  is  made  by  heating  the  carbonate.  It  unite, 
with  water  to  form  strontium  hydroxide  (Sr(OH)2),  which  is  used  in 
the  manufacture  of  beet  sugar.  Strontium  nitrate  (Sr(NO3)2)  and 
other  salts  of  strontium  color  a  flame  crimson,  and  are  widely  useo  i>-. 


Calcium,  Strontium,  and  Barium.  329 

making  fireworks,  especially  "red  fire."  The  latter  is  a  mixture  of 
potassium  chlorate,  shellac,  and  strontium  nitrate. 

The  production  of  the  crimson  colored  flame  is  the  test  for  stron- 
tium. 

Compounds  of  Barium.  — The  most  abundant  native  compounds  are 
witherite  (barium  carbonate,  BaCO3)  and  barite  (barium  sulphate, 
BaSO4) .  The  oxides,  BaO  and  BaO2,  have  already  been  mentioned  as 
a  source  of  oxygen.  Barium  hydroxide  (Ba(OH)2)  solution  is  often 
called  baryta  water,  and  it  forms  the  insoluble  barium  carbonate 
(BaCO3)  when  exposed  to  carbon  dioxide.  Barium  chloride  (BaCl2)  is 
used  in  the  laboratory  to  test  for  sulphuric  acid  and  soluble  sulphates, 
because  it  readily  interacts  with  them  and  forms  the  insoluble  barium 
sulphate  (BaSO4).  This  precipitated  salt  is  a  fine,  white  powder,  and 
being  cheap  and  heavy  it  is  a  common  adulterant  of  the  ordinary  white 
paint,  Ground  native  barium  sulphate  has  a  similar  use.  Barium  sul- 
phate is  also  used  to  increase  the  weight  of  paper  and  to  give  it  a  gloss. 
Barium  salts  color  a  flame  green,  and  barium  nitrate  (Ba(NO3)2)  is 
extensively  used  in  making  fireworks,  especially  "green  fire."  Com- 
mercial barium  sulphide  (BaS),  as  well  as  the  sulphides  of  calcium 
and  strontium,  shine  feebly  in  the  dark,  after  having  been  exposed  to  a 
bright  light.  On  account  of  this  property  they  are  used  in  making 
luminous  paint.  Soluble  barium  salts  are  poisonous. 

The  production  of  the  green  flame  is  the  test  for  barium. 

EXERCISES. 

1.  Name  the  alkaline  earth  metals.     What  is  the  symbol  of  each  ? 

2.  Name  several  compounds  of  calcium.     What  proportion  of  the 
earth's  crust  is  calcium  ? 

3.  Describe  the  preparation  and  state  the  properties  of  calcium. 

4.  What  is  the  formula  of  calcium  carbonate  ?     State  the  properties, 
occurrence,  and  uses  of  (a)  limestone,  and  (£)  marble. 

5.  State  the  essential  characteristics  of  (a)  calcite,  (£)  Iceland  spar, 
(c)  stalactites,    (d)  Mexican    onyx,    (e)  travertine,    (/)  coquina,  (g) 
chalk,  (Ji)  coral. 

6.  Review  the  properties  of  calcium  carbonate,  especially  its  solu- 
bility (see  Chapter  XIV). 

7.  State  the  uses  of  (a)  limestone,  (£)  marble,  (c)  chalk. 

8.  Describe  the  formation  of  (a)  limestone  caves,  (£)   chalk,  (<:) 
coral. 


jjo  Descriptive  Chemistry. 

9.  What  is  the  formula  and  chemical  name  of  lime  ?  State  the 
properties  and  uses  of  lime.  How  is  it  made  ?  State  the  equation  for 
the  chemical  change. 

10.  What  is  (a)  quicklime,  (fr)  slaked  lime,  (c)  hydraulic  lime,  (d) 
Portland  cement,  (e)  "  air-slaked  "  lime  ? 

1 1 .  What  is  the  formula  of  calcium  hydroxide  ?     How  is  it  formed  ? 
What  are  its  properties  ?     How  does  it  interact  with  carbon  dioxide  ? 
State  the  equation  for  the  reaction. 

12.  What  is  («)  limewater,  (b)  milk  of  lime,  (c~)  whitewash  ? 

13.  What  is  mortar  ?     How  is  it  prepared  ?     For  what  is  it  used  ? 
How  does  it  change  chemically  with  age  ?     What  is  plaster  ? 

14.  Discuss  the  occurrence  of  calcium  sulphate.      State  the  chief 
properties  of  (a}  gypsam,  (b}  selenite,  (c)  alabaster,  (d)  satin  spar. 
For  what  are  gypsum  and  alabaster  used  ? 

15.  What  is  plaster  of  Paris  ?    Why  so  called  ?     How  is  it  pre- 
pared ?     What  is  its  chief  property  ?    What  are  its  uses  ?     What  is 
the  chemical  explanation  of  "  setting ''  ?     What  is  stugco  ? 

1 6.  What  is  hard  water  ?     How  does  it  act  with  soap  ?     What  is 
(#)  temporary  hardness,  and  (b)  permanent  hardness  ?     How  may  each 
be  removed  ?    What  is  soft  water  ?    Why  is  rain  water  often  called  soft 
water  ? 

17.  Summarize  the  properties   of  calcium  chloride.     What  is   its 
formula  ?     How  is  it  prepared  ? 

1 8.  Review  the  essential  properties  of  (a)  calcium  fluoride,  (b)  cal- 
cium carbide,  (c)  tricalcium  phosphate,  (d}  bleaching  powder. 

19.  What  is  the  test  for  (a)  calcium,  (b)  strontium,  (c)  barium  ? 

20.  State  the  use  of  (a)  strontium  hydroxide,  and  {b)  strontium 
nitrate. 

21.  For  what  are   (a)   barium  hydroxide,   (b}  barium  nitrate,  (V) 
barium  sulphide,  and  (//)  barium  chloride  used  ?      Describe  barium 
sulphate. 

PROBLEMS. 

1.  What   is   the  per  cent   of  calcium   in    (a)  marble    (CaCO<3), 
(£)  gypsum  (CaSO4 .  2  H.,O),   (V)  fluor  spar  (CaF2),   (d}  superphos- 
phate of  lime  (CaH4(PO4)2)  ? 

2.  How  many  tons  of  limestone  must  be  heated  to  produce  100 
tons  of  quicklime  ?     (Assume  CaCO3  =  CaO  +  CO2.) 

3.  Calculate  the  simplest  formula  of  a  compound  having  the  per- 
centage composition  Ca  =  40,  C  =  12,  O  =  48. 


CHAPTER  XXIV. 
MAGNESIUM,  ZINC,  CADMIUM,   AND  MERCURY. 

THESE  elements  form  a  natural  group,  though  the  mem- 
bers are  not  so  closely  related  as  the  alkali  and  alkaline 
earth  groups.  Zinc  and  cadmium  are  much  alike,  and  both 
also  resemble  magnesium.  Mercury  differs  somewhat 
from  zinc  and  cadmium,  but  resembles  copper. 

MAGNESIUM. 

Occurrence  of  Magnesium.  —  Magnesium  is  never 
found  free.  In  combination  it  is  widely  distributed  and 
very  abundant,  constituting  about  2.5  per  cent  of  the 
earth's  crust.  Dolomite  is  magnesium  calcium  carbonate 
(CaMg(CO3)2) ;  it  forms  whole  mountain  ranges  and  vast 
deposits;  beds  hundreds  of  feet  thick  cover  thousands  of 
square  miles  in  the  upper  Mississippi  valley.  Dolomite 
closely  resembles  marble  and  limestone.  Magnesium 
carbonate  is  also  abundant.  Many  of  the  Stassfurt 
salts  contain  magnesium,  for  example,  kainite  (KC1, 
MgSO4  .  3  H2O),  carnallite  (KC1,  MgCl2.6  H2O),  and 
kieserite  (MgSO4 .  H2O).  It  is  also  a  component  of 
serpentine,  talc,  soapstone,  asbestos,  meerschaum,  and 
other  silicates.  The  sulphate  and  chloride  are  found  in 
sea  water  and  in  mineral  springs. 

Through  the  decay  of  rocks,  magnesium  compounds  find  their  way 
into  the  soil,  from  which  they  are  taken  up  by  plants.  Magnesium 
phosphates  are  found  in  the  bones  of  animals  and  the  seeds  of  grains, 
and  also  in  guano, 


Descriptive  Chemistry. 


D 


II      (7= 


Preparation  of  Magnesium.  —  Magnesium  was  formerly  prepared  by 
reducing  the  chloride  with  sodium.  It  is  now  economically  manufac- 
tured by  electrolysis.  A  sketch  of  the  essential  parts  of  the  apparatus 
is  shown  in  Figure  67.  Carnallite  is  put  into  the  cylindrical  iron  vessel, 
C,  which  is  the  cathode.  This  is  closed  by  the  air-tight  cover  through 
which  pass  the  pipes,  Z>,  D ',  for  conveying  inert  gases  into  and  out  of 
the  apparatus.  The  carbon  anode,  A,  dips  into  the  carnallite  and 
is  inclosed  by  the  porcelain  cylinder,  B,  which  is  provided  with  a 

pipe,  E,  for  the  escape  of  the  chlorine 
liberated  at  the  anode.  The  carnallite 
is  kept  fused  by  external  heat.  When 
the  current  passes,  the  chlorine  liberated 
at  the  anode  escapes  through  E,  and 
the  magnesium  liberated  at  the  cathode 
floats  on  the  fused  carnallite  and  is  pre- 
vented from  oxidizing  by  the  inert  gas 
supplied  through  D.  The  porcelain 
cylinder,  B,  prevents  the  chlorine  from 

escaping  into  the  larger  vessel.      The 
FIG.    67.  —  Apparatus  'for    the 

manufacture  of  magnesium  by  the     molten  magnesium  is  carefully  removed 
electrolysis  of  carnallite.  at  intervals. 

Properties  of  Magnesium.  —  Magnesium  is  a  lustrous, 
silvery  white  metal.  It  is  a  light  metal,  the  specific  grav- 
ity being  only  1.75.  It  is  tenacious  and  ductile,  and  when 
hot  may  be  drawn  into  wire  or  pressed  into  ribbon,  the 
latter  being  a  common  commercial  form.  It  melts  at  a  red 
heat  and  may  be  cast  into  different  shapes.  At  a  high 
temperature  it  volatilizes.  It  is  easily  kindled  by  a  match 
or  candle,  and  burns  with  a  dazzling  white  light,  producing 
dense  white  clouds  of  magnesium  oxide  (MgO).  It  does 
not  tarnish  in  dry  air,  but  in  moist  air  it  is  soon  covered 
with  a  film  of  oxide.  It  liberates  hydrogen  from  acids. 
Heated  in  nitrogen,  it  forms  magnesium  nitride  (Mg3N2, 
see  Composition  of  Ammonia). 

Uses  of  Magnesium.  —  Magnesium  in  the  form  of  pow- 
der is  used  chiefly  in  taking  flash-light  photographs. 


Magnesium,  Zinc,  Cadmium,  and  Mercury.    333 

Small  quantities  are  used  in  making  fire-works ;  and 
both  the  powder  and  wire  are  used  in  the  chemical 
laboratory. 

Magnesium  Oxide,  MgO,  is  a  white,  bulky  powder.  It 
is  formed  when  magnesium  burns  in  the  air,  but  it  is  man- 
ufactured by  gently  heating  magnesium  carbonate,  just  as 
lime  is  made  from  limestone.  It  is  often  called  magnesia, 
or  calcined  magnesia.  The  native  oxide  is  the  mineral 
periclase.  Magnesia  dissolves  with  difficulty  in  water, 
forming  magnesium  hydroxide  (Mg(OH)2).  A  mixture 
of  magnesia  and  water,  with  or  without  magnesium  chlo- 
ride, hardens  on  exposure  to  the  air,  and  is  often  used  as  a 
cement  or  artificial  stone.  Native  magnesium  hydroxide 
is  the  mineral  brucite.  Like  lime,  magnesia  withstands 
a  high  temperature,  and  is,  therefore,  used  as  the  chief 
ingredient  of  a  protective  mixture  for  steam  pipes  and  ves- 
sels which  are  subjected  to  great  heat.  Magnesia  is  used 
as  a  medicine  for  dyspepsia  and  an  antidote  for  poisoning 
by  mineral  acids. 

Magnesium  Sulphate,  MgSO4,  is  a  white  solid.  There 
are  several  crystallized  varieties.  The  native  salt  kie- 
serite  (MgSO4 .  H2O)  when  added  to  water  changes  into 
Epsom  salts  (MgSO4 .  ;H2O).  This  variety  was  first 
found  in  the  mineral  spring  at  Epsom,  England.  It  is 
very  soluble  in  water,  and  its  solution  has  a  bitter  taste. 
It  is  extensively  used  as  a  medicine,  in  manufacturing  sul- 
phates of  sodium  and  potassium,  as  a  fertilizer  in  place  of 
gypsum,  and  as  a  coating  for  cotton  cloth. 

Magnesium  Chloride,  MgCl2,  is  a  white  solid.  It  is  a  by- 
product in  the  preparation  of  potassium  chloride.  The  crystallized 
salt  (MgCl2 .  6  H.,O)  is  very  deliquescent.  Magnesia  mixture  is  a  mix- 
ture of  magnesium  chloride,  ammonium  chloride,  and  ammonium  hy- 
droxide ;  it  is  used  in  chemical  analysis. 


334  Descriptive  Chemistry. 

Magnesium  Carbonate,  MgCO3,  occurs  native  as  magnesite,  and 
combined  with  calcium  carbonate  as  dolomite.  The  commercial  salt 
known  as  magnesia  alba,  or  simply  magnesia,  is  a  complex  compound 
(Mg(OH)2,  4  MgCO3  •  4  H2O) .  Several  of  these  complex  basic  carbon- 
ates are  known.  Many  face  powders  consist  chiefly  of  magnesia  alba. 

It  was  during  an  investigation  of  magnesia  alba  that  Black  discov- 
ered carbon  dioxide  and  showed  the  close  relation  between  analogous 
compounds  of  magnesium  and  calcium. 

Miscellaneous.  —  Besides  the  oxide  and  sulphate,  other  compounds 
are  used  as  medicines.  Fluid  magnesia,  prepared  by  dissolving  mag- 
nesium carbonate  in  water  containing  carbon  dioxide,  is  a  mild  laxative. 
Magnesium  citrate  has  a  similar  action ;  it  is  an  effervescing  mixture 
prepared  from  sodium  bicarbonate,  tartaric  and  citric  acids,  sugar,  and 
magnesium  sulphate. 

ZINC. 

Occurrence  of  Zinc.  —  Free  zinc  is  never  found.  The 
ores  of  zinc  are  not  numerous,  but  are  widely  distributed. 
The  chief  ores  are  zinc  sulphide  (sphalerite,  zinc  blende, 
ZnS),  zinc  carbonate  (smithsonite,  ZnCO3),  zinc  silicate 
(calamine,  H2Zn2SiO5),  and  red  zinc  oxide  (zincite,  ZnO). 
Franklinite  and  willemite  are  ores  of  zinc  containing 
manganese  and  iron.  Gahnite  has  the  composition 
ZnAl2O4. 

Zinc  ores  are  found  in  Germany,  Italy,  France,  Greece,  Spain,  Austria- 
Hungary,  Belgium,  England,  and  the  United  States.  Missouri  and 
Kansas  contain  large  deposits  of  the  sulphide,  while  the  other  ores 
occur  chiefly  in  New  Jersey.  About  143,000  tons  of  zinc  were  pro- 
duced in  the  United  States  in  1902,  and  over  60  per  cent  came  from 
Missouri-Kansas.  This  was  the  largest  amount  ever  produced  in  a 
single  year. 

Metallurgy  of  Zinc.  —  Zinc  is  easily  smelted.  The  ores 
are  first  roasted  to  change  them  into  the  oxide,  thus  — 

ZnCO3      =       ZnO     +         CO2 

Zinc  Carbonate        Zinc  Oxide      Carbon  Dioxide 


Magnesium,  Zinc,  Cadmium,  and  Mercury.    335 

ZnS      +30=     ZnO     +         SO2 
Zinc  Sulphide         Oxygen      Zinc  Oxide     Sulphur  Dioxide 

The  oxide  is  then  reduced  by  heating  it  with  charcoal. 
This  operation  is  conducted  in  earthenware  tubes  or  fire- 
clay crucibles  connected  with  iron  receivers  into  which  the 
zinc  vapor  passes ;  at  first  it  condenses  as  a  powder  known 
as  zinc  dust,  somewhat  as  sulphur  forms  flowers  of  sul- 
phur ;  but  it  finally  condenses  as  a  liquid,  which  is  drawn 
off  at  intervals  and  cast  into  bars  or  plates.  The  impure 
zinc  thus  obtained  is  called  spelter ;  it  is  freed  from  carbon, 
lead,  iron,  cadmium,  and  arsenic  by  repeated  distillation, 
often  under  reduced  pressure. 

Properties  of  Zinc.  —  Zinc  is  a  bluish  white,  lustrous 
metal.  Its  physical  properties  vary  with  the  temperature. 
At  ordinary  temperatures  it  is  brittle,  but  at  100°  —  150°  C. 
it  is  soft  and  may  be  rolled  into  sheets  and  drawn  into 
wire,  while  its  specific  gravity  rises  from  6.9  to  7.2.  Zinc 
which  has  been  rolled  or  drawn  does  not  become  brittle 
upon  cooling.  At  200°  C.  it  again  becomes  brittle  and 
can  be  easily  pulverized.  It  melts  at  about  433°  C.  and 
boils  at  about  940°  C.  Heated  in  the  air  above  its  melting 
point,  zinc  burns  with  a  bluish  green  flame,  forming  white 
zinc  oxide  (ZnO).  Zinc  does  not  tarnish  in  dry  air,  but 
ordinarily  it  becomes  coated  with  a  dark  film.  Commercial 
vzinc  interacts  with  acids  and  usually  liberates  hydrogen. 
With  hot  solutions  of  sodium  and  potassium  hydroxides,  it 
forms  zincates  and  liberates  hydrogen,  thus  — 

2KOH          +  Zn    =       H2      +        K2ZnO2. 
Potassium  Hydroxide       Zinc          Hydrogen      Potassium  Zincate 

Pure  zinc  interacts  with  acids  if  in  contact  with  a  platinum 
wire,  or  if  copper  sulphate  solution  is  added.  Like  copper, 


33 6  Descriptive  Chemistry. 

zinc  withdraws  other  metals  (e.g.  lead  and  mercury)  from 
their  solutions. 

The  vapor  density  of  zinc  requires  the  molecular  weight  67.6.  Since 
the  atomic  weight  is  65.4,  a  molecule  of  the  vapor  contains  only  one 
atom. 

Uses  of  Zinc.  —  Zinc  in  stick  or  plates  is  extensively 
used  as  the  positive  plate  in  electric  batteries.  Sheet  zinc 
is  used  as  a  lining  for  tanks,  and  as  the  protective  cover- 
ing which  is  placed  behind  and  beneath  stoves.  Iron 
dipped  into  melted  zinc  becomes  coated  with  zinc  and  is 
called  galvanized  iron ;  it  does  not  rust  easily  and  is  widely 
used  for  roofs,  pipes,  cornices,  and  water  tanks.  Telegraph 
wire  is  also  galvanized.  Zinc  dust  is  used  in  the  cyanide 
process  of  extracting  gold  and  in  many  chemical  experi- 
ments in  the  laboratory.  Brass,  German  silver,  and  other 
alloys  contain  zinc  (see  Alloys  of  Copper).  Antifriction 
metals,  which  are  used  for  bearings,  are  alloys  of  zinc. 
Babbitt's  metal,  for  example,  contains  69  per  cent  of  zinc, 
19  of  tin,  4  of  copper,  3  of  antimony,  and  5  of  lead. 

Compounds  of  Zinc.  —  Native  zinc  oxide  is  red,  owing 
to  the  presence  of  manganese,  but  the  pure  oxide  is  white 
when  cold  and  yellow  when  hot.  It  is  formed  when  zinc 
burns,  and  is  manufactured  in  this  way  or  by  heating  zinc 
carbonate.  It  is  often  called  "zinc  white"  or  "  Chinese 
white,"  and  is  used  to  make  a  white  paint  which  is  not  dis- 
colored by  the  atmosphere.  Native  zinc  sulphide  is  yel- 
low, brown,  or  black  on  account  of  impurities,  but  the  pure 
sulphide  is  white.  The  latter  is  formed  as  a  jelly  like  pre- 
cipitate when  hydrogen  sulphide  is  passed  into  an  alkaline 
solution  of  a  zinc  salt ;  it  is  decomposed  by  a  mineral  acid. 
Zinc  sulphide  is  also  used  as  a  white  pigment.  Zinc 
sulphate  is.  formed  by  the  interaction  of  zinc  and  dilute 
sulphuric  acid.  Large  quantities  are  made  by  roasting 


Magnesium,  Zinc,  Cadmium,  and   Mercury.    337 

the  sulphide  in  a  limited  supply  of  oxygen  and  extracting 
the  sulphate  with  water.  It  is  a  white,  crystallized  solid 
(ZnSO4 .  7  H2O),  which  effloresces  in  the  air,  and  when 
heated  to  100°  C.  loses  most  of  its  water  of  crystallization. 
The  crystallized  salt  is  called  white  vitriol.  It  is  used  in 
dyeing  and  calico  printing,  as  a  disinfectant,  and  as  a  medi- 
cine. It  is  poisonous,  but  can  be  safely  used  externally  to 
relieve  inflammation.  Zinc  chloride  (ZnCl2)  is  a  white, 
deliquescent  solid,  prepared  by  dissolving  zinc  in  hydro- 
chloric acid  and  evaporating  the  solution  until  a  sample 
solidifies  on  cooling.  It  is  used  in  surgery,  and  also  as  a 
constituent  of  a  mixture  for  filling  teeth ;  large  quantities 
are  used  to  preserve  wood,  especially  railroad  ties,  from 
decay,  nearly  1500  tons  being  annually  consumed  for  this 
purpose.  Zinc  hydroxide  (Zn(OH)2)  is  formed  by  the 
interaction  of  sodium  or  potassium  hydroxide  and  the  solu- 
tion of  a  zinc  salt.  An  excess  of  the  alkaline  hydroxide 
changes  the  zinc  hydroxide  into  a  zincate. 

Tests  for  Zinc.  —  The  formation  of  the  sulphide  or  hydroxide,  as 
above  described,  serves  as  the  test  for  zinc.  A  green  incrustation  is 
produced  when  zinc  compounds  are  heated  on  charcoal  and  then  mois- 
tened with  a  cobaltous  nitrate  solution. 

Cadmium,  Cd,  is  an  uncommon  metal,  frequently  found  in  zinc  ores. 
It  occurs  native  as  a  sulphide  (greenockite,  CdS).  It  is  white,  lustrous, 
and  rather  soft.  Its  specific  gravity  is  8.6,  and  its  melting  point  is 
about  320°  C.  Cadmium  is  a  constituent  of  certain  fusible  alloys  (see 
Bismuth).  Wood's  metal  contains  12  per  cent  of  cadmium.  The  most 
important  compound  is  cadmium  sulphide  (CdS).  This  is  a  bright 
yellow  solid,  formed  by  adding  hydrogen  sulphide  to  the  solution  of  a 
cadmium  compound.  It  is  used  as  an  artist's  color.  Its  formation  also 
serves  as  the  test  for  cadmium. 

MERCURY. 

Occurrence  of  Mercury.  —  Native  mercury  is  occasion- 
ally found  in  minute  globules,  but  the  most  abundant  ore 


jj  8  Descriptive  Chemistry. 

is  mercuric  sulphide  (cinnabar,  HgS).  The  ore  is  mined 
in  Spain,  Austria,  Russia,  Italy,  and  Mexico ;  in  the  United 
States  large  quantities  are  obtained  in  California,  and 
deposits  were  recently  opened  in  Texas. 

The  annual  production  of  the  United  States  for  several  years  has 
been  about  1000  tons. 

Mercury  has  been  known  for  ages  as  quicksilver.  The  Latin  name, 
hydrargyrum,  which  gives  us  the  symbol  Hg,  means  literally  "  water 
silver,"  emphasizing  the  fact,  so  well  known,  that  mercury  looks  like 
silver  and  flows  like  water. 

Preparation  of  Mercury.  —  Mercury  is  readily  prepared 
by  roasting  cinnabar  in  a  current  of  air.  Sulphur  dioxide 
and  mercury  are  formed,  thus  — 

HgS         +         02  Hg         +         S02 

Cinnabar  Oxygen  Mercury  Sulphur  Dioxide 

The  sulphur  dioxide  is  usually  allowed  to  escape,  but  the 
mercury  vapor  is  condensed  by  passing  it  into  large  cham- 
bers, or  through  pear-shaped  retorts  or  pipes,  called  aludels 
(see  Iodine).  Crude  mercury  is  freed  from  dirt  and  me- 
chanical impurities  by  pressing  it  through  linen  or  chamois 
leather,  but  it  must  be  distilled  to  separate  it  from  dissolved 
metals,  such  as  lead  or  zinc.  It  can  also  be  purified  by 
treatment  with  dilute  nitric  acid.  Mercury  is  sent  into 
commerce  in  strong  iron  flasks,  holding  about  75  pounds. 
Properties  of  Mercury.  —  Mercury  is  a  bright,  silvery 
metal,  and  is  the  only  one  which  is  liquid  at  ordinary  tem- 
peratures. It  solidifies  at  about  —  39.5°  C.  It  is  a  heavy 
metal,  the  specific  gravity  being  13.59.  It  is  slightly  vola- 
tile even  at  ordinary  temperatures,  and  the  vapor  is  poison- 
ous. Mercury  does  not  tarnish  in  the  air,  unless  sulphur 
compounds  are  present.  At  a  high  temperature,  it  com- 
bines slowly  with  oxygen  to  form  the  red  oxide  (HgO). 


Magnesium,  Zinc,  Cadmium,  and  Mercury.     339 

Hydrochloric  acid  and  cold  sulphuric  acid  do  not  affect  it ; 
hot  concentrated  sulphuric  acid  oxidizes  it,  and  nitric  acid 
changes  it  into  nitrates. 

The  vapor  density  of  mercury  requires  the  molecular  weight  198.72. 
Since  the  atomic  weight  is  200,  a  molecule  of  the  vapor  contains  only 
one  atom. 

Amalgams  are  alloys  of  mercury  with  other  metals. 
They  are  easily  prepared  by  mixing  the  constituents. 
Sometimes  the  union  is  violent  as  in  the  preparation  of 
sodium  amalgam.  Amalgamated  zinc  is  usually  used  in 
electric  batteries  to  prevent  unnecessary  loss  of  the  zinc. 
Tin  amalgam  is  sometimes  used  to  coat  mirrors.  Amal- 
gams of  certain  metals  are  used  as  a  filling  for  teeth.  Care 
should  be  taken,  while  handling  mercury,  not  to  let  it  come 
in  contact  with  rings  or  jewelry,  since  gold  amalgam  is 
readily  formed. 

Uses  of  Mercury.  —  Mercury  is  used  in  making  ther- 
mometers, barometers,  and  some  kinds  of  air  pumps.  Its 
extensive  use  in  extracting  gold  and  silver  has  been  men- 
tioned (see  Amalgamation).  Large  quantities  are  used  in 
preparing  certain  medicines  and  explosives  (e.g.  fulminating 
mercury,  which  is  used  in  cartridges). 

Compounds  of  Mercury.  —  Mercury,  like  copper,  forms  two  classes  of 
compounds  —  the  mercurous  and  the  mercuric.  Mercuric  oxide  (HgO) 
is  a  red  powder,  produced  by  heating  mercury  in  air  or  by  heating  a 
mixture  of  mercury  and  mercuric  nitrate.  As  we  have  already  seen, 
mercuric  oxide  is  decomposed  by  heat  into  mercury  and  oxygen.  A 
yellow  variety  is  produced  by  the  interaction  of  sodium  hydroxide  and  a 
mercuric  salt,  thus  — 

2NaOH     +     Hg(NO3)2     =     HgO     +     2  NaNO3     +     H2O 
Sodium  Mercuric  Mercuric          Sodium 

Hydroxide  Nitrate  Oxide  Nitrate 

Mercurous  chloride  (Hg2Cl2  or  HgCl)  is  a  white,  tasteless  powder, 
insoluble  in  water.  It  is  formed  when  a  chloride  and  mercurous  nitrate 


34-O  Descriptive  Chemistry. 

interact,  but  it  is  manufactured  by  heating  a  mixture  of  mercuric  chloride 
and  mercury.  Under  the  name  of  calomel  it  is  extensively  used  as  a 
medicine.  Mercuric  chloride  (HgCl2)  is  a  white,  crystalline  solid,  solu- 
ble in  water  and  in  alcohol.  It  is  prepared  by  heating  a  mixture  of 
mercuric  sulphate  and  common  salt.  It  is  a  violent  poison.  The  best 
antidote  is  the  white  of  a  raw  egg.  The  albumen  forms  an  insoluble 
mass  with  the  poison,  which  may  then  be  removed  mechanically  from 
the  stomach.  The  common  name  of  mercuric  chloride  is  corrosive 
sublimate.  It  has  strong  antiseptic  properties,  and  is  extensively  used 
in  surgery  to  protect  wounds  from  the  harmful  action  of  germs ;  taxi- 
dermists sometimes  use  it  to  preserve -skins,  and  it  has  many  serviceable 
applications  as  a  medicine  and  disinfectant.  It  is  usually  used  as  a 
dilute  solution  (i  part  to  1000  parts  of  water).  Native  mercuric  sul- 
phide or  cinnabar  (HgS)  is  a  red,  crystalline  solid.  When  hydrogen 
sulphide  is  passed  into  a  solution  of  a  mercuric  salt,  mercuric  sulphide 
is  formed  as  a  black  powder;  this  variety,  when  heated,  changes  into 
red  crystals. 

Vermilion  is  artificial  mercuric  sulphide.  It  is  manufactured  either 
(i)  by  grinding  together  mercury  and  sulphur,  and  treating  this  mass 
with  caustic  potash  solution,  or  (2)  by  heating  mercury  and  sulphur  in 
iron  pans  and  subliming  the  black  mass.  In  both  processes  the  product 
must  be  carefully  ground,  washed,  and  dried.  Chinese  vermilion  is  the 
best  quality.  Vermilion  has  a  brilliant  red  color,  and,  although  expen- 
sive, is  widely  used  to  make  red  paint. 

Mercurous  Nitrate  (HgNO3  or  Hg2(NO3)2)  and  mercuric  nitrate 
(Hg(NO3)2)  are  prepared  by  treating  mercury  respectively  with  cold 
dilute  nitric  acid,  and  with  hot  concentrated  nitric  acid.  They  are 
white,  crystalline  solids. 

EXERCISES. 

1.  Name  the  chief  native  compounds  of  magnesium.     What  pro- 
portion of  the  earth's  crust  is  magnesium  ? 

2.  Describe  the  manufacture  of  magnesium  by  the  electrolysis  of 
carnallite. 

3.  Summarize  the  properties  of  magnesium.     State  its  uses. 

4.  What  is  the  formula  and  chemical  name  of  magnesium  ?     How 
is  magnesia  formed  ?     State  its  properties  and  uses. 

5.  Describe  the  different  varieties  of  magnesium  sulphate.     State 
the  uses  of  Epsom  salts. 


Magnesium,  Zinc,  Cadmium,  and  Mercury.    341 

6.  What  is  the  formula  of  magnesium  carbonate  ?     What  is  (#) 
magnesite,  (£)  dolomite,  (c)  magnesia  alba?     For  what  is  the  last  sub- 
stance used  ? 

7.  Name  the  chief  ores  of  zinc.     Discuss  their  occurrence. 

8.  Describe  the  metallurgy  of  zinc.     What  is  (a)  zinc  dust,  and 
(£)  spelter  ?     How  is  zinc  purified  ? 

9.  Summarize  (a)  the  physical  properties  of  zinc,  and  (£)  the  chem- 
ical properties. 

10.  State  the  uses  of  zinc. 

11.  Review  the  alloys  of  copper  which  also  contain  zinc.     What 
alloys  are  largely  zinc  ? 

12.  Describe  native  and  pure  zinc  oxide.     For  what  is  the  latter 
used  ? 

13.  Describe  zinc  sulphate.      How  is  it  formed  and  for  what  is  it 
used? 

14.  Describe  zinc  chloride.     For  what  is  it  used? 

15.  What  are  the  tests  for  zinc  ? 

1 6.  State  the  properties  and  uses  of  (a)  cadmium,  and  (£)  cadmium 
sulphide. 

17.  What  is  the  chief  ore  of  mercury  ?     Where  is  it  found  ? 

1 8.  What  is  the  symbol  of  mercury?     What  is  the  literal  meaning 
of  the  word  from  which  it  is  formed  ? 

19.  Describe  the  preparation  and  purification  of  mercury.     How  is  it 
transported  ? 

20.  Summarize  the  properties  of  mercury. 

21.  What  are  amalgams  ?     Name  three,  and  state  the  use  of  each. 

22.  For  what  is  mercury  used  ? 

23.  Describe  mercuric  oxide.     What  historical  interest  has  it  ? 

24.  Describe  mercurous  chloride.     What  is  its  commercial  name? 
State  its  use. 

25.  Describe   mercuric  chloride.     What   is  its  commercial   name? 
How  does  it  differ  from  mercurous  chloride  ?     State  its  use. 

26.  What  is  the  formula  and  chemical  name  of  cinnabar  ?     Describe 
cinnabar.     What  is  vermilion  ?     How  is  it  manufactured?     State  its 
use. 

27.  What  is  (a)  magnesia,   (£)  Epsom   salts,  (c)  galvanized  iron, 
(d)  Chinese  white,  0)  white  vitriol,  (/)  calomel,  (g)  corrosive  subli- 
mate ? 


342  Descriptive  Chemistry. 


PROBLEMS. 

1.  How  much    magnesium  will  be  formed  by  heating  100  gm.  of 
potassium     with     magnesium     chloride  ?       (Assume     K2  +  MgCl2  = 
Mg  +  2  KC1.) 

2.  What  is  the  per  cent  of  magnesium  in  (#)  magnesite  (MgCO3), 
(£)  dolomite  (MgCa(CO3)2),  (c}  Epsom  salts  (MgSO4  •  7  H,O  )  ? 

3.  What  is  the  per  cent  of  zinc  in  (a)  zinc  sulphate  (ZnSO4),  (b} 
zinc  sulphide  (ZnS),  (c)  zinc  chloride  (ZnCl2),  (d)  zinc  oxide  (ZnO)  ? 

4.  How  much  zinc  sulphate  can  be  prepared  from  65  gm.  of  zinc  ? 
From  130  gm.?     From  720  gm.? 

5.  How  much  mercury  is  formed  by  decomposing  400  gm.  of  cin- 
nabar ?     (Assume  HgS  +  O2  =  Hg  +  SO2.) 

6.  What  is  the  per  cent  of  mercury  in  (a)  mercuric  oxide  (HgO), 
(b)  calomel  (Hg2Cl2),  (c)  corrosive  sublimate  (HgCl2)  ? 


CHAPTER  XXV. 
ALUMINIUM. 

Occurrence.  —  Aluminium  does  not  occur  free  in  nature, 
but  its  compounds  are  numerous,  abundant,  and  widely 
distributed.  About  8  per  cent  of  the  earth's  crust  is 
aluminium;  it  is,  therefore,  the  most  abundant  metal. 
Many  common  rocks  and  minerals  are  silicates  of  alumin- 
ium and  other  metals,  e.g.  feldspar  and  mica,  which  make 
up  a  large  part  of  granite  and  gneiss.  Clay  and  slate  are 
mainly  silicate  of  aluminium,  formed  by  the  decomposition 
of  complex  aluminium  minerals.  Corundum  and  emery 
are  aluminium  oxide  (A12O3)  more  or  less  impure.  Baux- 
ite is  an  hydroxide  of  aluminium  (H4A12O5).  Cryolite  is  a 
fluoride  of  aluminium  and  sodium  (Na3AlF6). 

Aluminium  was  first  obtained  as  a  fine  powder  by  Wohler  in  1827. 
Deville,  in  1854,  prepared  it  in  compact  form  and  laid  the  foundation 
of  the  industry  which  is  being  developed  by  Hall. 

Davy  proposed  the  name  alumium,  i.e.  alum  +  him,  to  emphasize  the 
relation  of  the  metal  to  the  well-known  substance,  alum.  The  word 
alumium  was  changed  first  to  aluminum  and  then  to  aluminium.  Some 
authorities  derive  the  word  alumium  from  the  Latin  word  alumen,  or 
from  alumina,  the  common  name  of  aluminium  oxide. 

Metallurgy.  —  Aluminium  is  obtained  from  its  oxide 
(A12O3)  by  electrolysis.  In  the  Hall  process,  which  is 
typical,  an  open,  iron  box  lined  with  carbon  is  made  the 
cathode  (Fig.  68).  The  anode  consists  of  carbon  bars 
hung  from  a  copper  rod,  which  can  be  lowered  as  the  car- 

343 


344 


Descriptive  Chemistry. 


bon  is  consumed.  The  process  is  essentially  as  follows : 
the  bottom  of  the  box  is  covered  with  cryolite,  the  anodes 
are  lowered,  and  the  box  is  then  filled  with  cryolite.  The 
current  is  turned  on,  and  in  its  resisted  passage  through 
the  cryolite  enough  heat  is  generated  to  melt  the  cryolite. 
Pure,  dry  aluminium  oxide  is  now  added.  This  is  decom- 
R  .  posed  into  aluminium 

and  oxygen.  The  oxy- 
gen unites  with  the 
carbon  of  the  anodes, 
forming  carbon  mo- 
noxide, which  burns  or 


escapes.      The   molten 

FlG.  68.  —  Apparatus  for  the  manufacture  of     aluminium   falls    to    the 
aluminium    by  the    electrolysis    of   aluminium 
oxide.     C  C  C  is  the  iron  box  which  serves  as 


the  cathode.      A,  A,  etc.   are   carbon   anodes 
attached  to  the  copper  rod,  R. 


bottom.  The  process 
is  continuous,  fresh 
aluminium  oxide  being 
added  and  the  molten  aluminium  being  drawn  off  at  inter- 
vals. The  cryolite  is  unchanged,  and  merely  acts  as  a 
solvent  for  the  aluminium  oxide. 

The  United  States  produced  about  7,000,000  pounds  of  aluminium 
in  1902,  and  the  output  is  annually  increasing.  This  was  all  produced 
at  Niagara  Falls.  In  the  Heroult  process,  which  is  used  in  Europe  and 
involves  essentially  the  same  principle  as  Hall's  process,  the  aluminium 
is  produced  as  an  alloy  (usually  of  copper) . 

Aluminium  was  prepared  until  about  1885  by  a  complicated  process, 
(i)  Bauxite  was  changed  into  aluminium  oxide  free  from  iron  by  fusion 
with  sodium  carbonate  and  treatment  with  carbon  dioxide.  (2)  The 
aluminium  oxide  was  then  changed  into  aluminium  sodium  chloride  by 
fusion  with  sodium  chloride  and  charcoal  and  subsequent  treatment  with 
chlorine.  (3)  This  chloride  was  reduced  by  sodium,  thus  — 

A1C13       +       3Na       =       Al       +       3  NaCl 
Aluminium         Sodium      Aluminium        Sodium 
Chloride  Chloride 


Aluminium.  345 

The  sodium  for  this  operation  was  prepared  by  the  Castner  process  (see 
Sodium),  and  the  two  industries  were  developed  simultaneously. 

The  extensive  application  of  the  electrolytic  method  has  reduced  the 
price  of  aluminium  from  about  $12  a  pound  during  1862-1887  to  about 
30  cents  in  1902. 

Properties. —  Aluminium  is  a  bluish  white  metal.  It  is 
very  light  compared  with  other  common  metals,  since  its 
specific  gravity  is  only  about  2.6 ;  this  value  is  one  third 
that  of  iron.  It  is  ductile  and  malleable,  and  is  often 
sold  in  the  form  of  wire  and  sheets ;  it  must  be  annealed 
frequently  during  the  hammering  or  drawing.  It  is  a 
good  conductor  of  heat  and  electricity.  Its  tensile 
strength  is  about  as  great  as  that  of  cast  iron.  It  melts  at 
about  660°  C.,  and  may  be  cast  and  welded,  but  not  readily 
soldered  so  as  to  produce  a  permanent  joint.  The  cap  of 
the  Washington  Monument  is  a  casting  of  aluminium 
which  weighs  about  eight  and  a  half  pounds.  Pure  alu- 
minium is  only  very  slightly  oxidized  by  air.  Hydrochlo- 
ric acid  changes  it  into  aluminium  chloride,  thus  — 

2A1-   +     6HC1     =     2A1C13     +     3H2 

Aluminium     Hydrochloric          Aluminium         Hydrogen 
Acid        .  Chloride 

Under  ordinary  conditions  nitric  and'  sulphuric  acids  do 
not  affect  it.  Sodium  and  potassium  hydroxides  change  it 
into  aluminates,  thus  — 

6NaOH      +     2A1     =     2  Na3AlO3     +     3  H2 
Sodium  Hydroxide     Aluminium     Sodium  Alumkiate       Hydrogen 

The  properties  of  aluminium  are  modified  by  the  presence  of  impuri- 
ties. The  usual  impurities  are  iron,  other  metals,  and  silicon.  Some 
of  these,  especially  the  iron  and  silicon,  come  from  the  raw  products 
used  in  its  manufacture.  They  tend  to  make  the  metal  harder  and  more 
active  chemically,  but  less  malleable,  ductile,  and  tenacious.  If  it  were 
not  for  the  presence  of  these  impurities  in  clay,  this  substance  would  be 
a  cheap  and  inexhaustible  source  of  aluminium. 


346  Descriptive  Chemistry. 

Uses.  —  The  varied  properties  of  aluminium  adapt  it  to 
numerous  uses.  It  is  made  into  the  metallic  parts  of  mili- 
tary outfits,  caps  for  fruit  jars,  surgical  instruments,  cook- 
ing utensils,  tubes,  the  framework  and  fittings  of  boats  and 
air  ships,  telephone  receivers,  scientific  apparatus,  parts  of 
opera  glasses  and  telescopes,  the  framework  of  cameras, 
stock  patterns  for  foundry  work,  and  hardware  samples. 
Its  attractive  appearance  has  led  to  its  extensive  use  as  an 
ornamental  metal,  both  in  interior  decorative  work  and  in 
numerous  small  objects,  such  as  trays,  picture  frames, 
hairpins,  and  combs.  Aluminium  leaf  is  used  for  decorat- 
ing book  covers  and  signs ;  the  powder  is  likewise  used  as 
a  protective  and  attractive  coating  for  letter  boxes,  steam 
pipes,  lamp-posts,  radiators,  smokestacks,  and  other  metal 
objects  exposed  to  heat  or  the  weather.  During  the  last 
few  years  aluminium  wire  has  come  into  use  as  a  conductor 
of  electricity.  Large  quantities  of  aluminium  are  used  to 
reduce  oxides,  to  make  iron  and  steel  more  fluid,  and  to 
produce  sounder  castings.  The  applications  of  aluminium 
are  constantly  increasing. 

Alloys.  —  The  alloy  of  aluminium  and  copper — aluminium  bronze  — 
has  been  been  described  (see  Alloys  of  Copper) .  Magnalium  is  a  recent 
alloy  containing  from  75  to  90  per  cent  of  aluminium,  the  rest  being 
magnesium. 

Aluminium  Oxide,  A12O3,  is  the  only  oxide  of  alumin- 
ium. It  is  often  called  alumina,  as  silicon  dioxide  is  called 
silica.  Its  native  forms,  corundum  and  emery,  are  found 
in  Massachusetts,  New  Jersey,  Georgia,  Pennsylvania, 
North  Carolina,  and  Canada ;  large  quantities  come  from 
Asia  Minor  and  the  islands  near  Greece.  Emery  is  ex- 
tremely hard,  and  is  used  in  various  forms — powder,  cloth, 
paper,  and  wheels  —  to  grind  and  polish  hard  metals,  plate 


Aluminium. 


347 


glass,  etc.  The  crystallized  varieties  of  aluminium  oxide 
are  usually  known  as  corundum,  and  the  transparent, 
colored  kinds  have  long  been  prized  as  gems  (see  below). 

Alumina  may  be  prepared  by  burning  the  metal  or  by  heating  its 
hydroxide.  Thus  prepared,  it  is  a  white  powder,  insoluble  in  water, 
but  soluble  in  zfcids  and  in  the  caustic  alkalies.  It  melts  in  the  oxyhy- 
drogen  flame,  and  in  the  electric  furnace.  Heating  lessens  its  chemical 
activity.  When  alumina  or  any  other  compound  of  aluminium  is  heated, 
then  cooled  and  moistened  with  cobaltous  nitrate  solution  and  heated 
again,  the  mass  turns  a  beautiful  blue  color.  This  is  a  test  for  alu- 
minium. 

Aluminium  is  both  basic  and  acid,  that  is,  with  acids  it  forms  salts, 
like  aluminium  chloride,  while  with  bases  it  forms  aluminates. 

Gems  containing  Aluminium.  —  Corundum  (A12O3)  has  long  been 
found  as  crystals  in  Ceylon,  Siam,  Burma,  and  other  places  in  the 
Orient.  The  color  is  due  to  traces  of  impurities,  usually  oxides  of 
metals.  The  sapphire  is  blue,  and  the  ruby  is  red.  The  Oriental 
topaz  is  yellow,  the  Oriental  amethyst  is  purple,  and  the  Oriental 
emerald  is  green.  Montana  furnishes  many  sapphires,  the  output  in 
1901  being  valued  at  $90,000.  These  gems  may  be  artificially  produced 
by  dissolving  alumina  in  a  fused  substance,  adding  an  oxide  to  secure 
the  desired  color,  and  then  allowing  the  alumina  to  crystallize.  Spinels 
are  complex  compounds  of  aluminium.  The  typical  or  ruby  spinel  is 
magnesium  aluminate  (MgAl2O4).  It  resembles  the  true  ruby.  Other 
spinels  differ  from  the  ruby  spinel  both  in  color  and  in  composition. 
Turquoise  is  a  complex  aluminium  phosphate  containing  traces  of  cop- 
per. It  has  a  beautiful  robinVegg-blue  color,  is  compact,  and  may  be 
worked  into  various  shapes.  Formerly  turquoise  came  almost  exclu- 
sively from  Persia,  but  now  New  Mexico  meets  all  demands.  Nearly 
$120,000  worth  of  turquoise  are  mined  annually  m  that  state.  Topaz 
is  a  complex  aluminium  silicate  containing  fluorine.  It  is  usually  pale 
yellow,  and  is  found  in  many  localities.  Emerald  is,  next  to  diamond 
and  ruby,  the  most  precious  gem.  It  is  an  aluminium  silicate  con- 
taining the  rare  element  beryllium.  The  finest  specimens  have  a  deep 
emerald-green  color  and  come  from  Colombia,  South  America.  Garnet 
is  a  complex  silicate  of  aluminium  and  another  metal,  especially  cal- 
cium, magnesium,  iron,  or  manganese.  The  kind  used  as  a  gem  has  a 
deep  red  color  and  is  rather  abundant. 


348  Descriptive  Chemistry. 

Aluminium  Hydroxide,  A1(OH)3,  is  a  white,  jelly  like 
solid  formed  by  adding  an  hydroxide  to  the  solution  of  an 
aluminium  salt,  thus  — 

AlClg     +   3  NH4OH   =    A1(OH)3  +  3  NH4C1 

Aluminium  Ammonium  Aluminium         Ammonium 

Chloride  Hydroxide  Hydroxide          'Chloride 

It  is  insoluble  in  water.  It  interacts  with  strong  acids 
and  with  alkalies  (except  ammonium  hydroxide),  forming 
respectively  aluminium  salts  and  aluminates.  Thus  — 

A1(OH)8  +  3  HC1  =  A1C18  4-  3  H2O 
Aluminium  Hydrochloric  Aluminium  Water 
Hydroxide  Acid  Chloride 

A1(OH)3  +   3  NaOH   =   Na3AlO3  +   3  H2O 

Sodium  Sodium 

Hydroxide  Aluminate 

Bauxite  is  a  native  aluminium  hydroxide,  though  it  contains  iron 
and  silicon.  It  resembles  clay  in  texture  and  color.  The  vast  deposits 
found  at  Baux,  in  southern  France,  furnish  most  of  the  raw  material  for 
the  manufacture  of  aluminium,  though  about  twenty  thousand  tons  are 
annually  obtained  from  our  Southern  states,  chiefly  from  Georgia. 

Aluminium  Sulphate,  A12(SO4)3.  18  H2O,  is  a  white, 
crystalline  solid.  The  commercial  salt  has  a  variable  com- 
position ;  and,  if  pure,  it  dissolves  readily  and  completely 
in  water.  It  is  extensively  used  in  dyeing  and  paper 
making,  and  in  preparing  other  aluminium  compounds. 

Aluminium  sulphate  is  prepared  from  pure  clay,  bauxite,  or  cryolite. 
If  clay  or  bauxite  is  heated  with  sulphuric  acid  and  then  allowed  to 
cool,  the  product  is  impure  aluminium  sulphate,  known  as  "  alum  cake," 
or,  if  much  iron  is  present,  as  "  alumino  ferric  cake.1'  It  is  used  to 
purify  sewage  and  for  other  purposes  where  iron  and  the  other  impuri- 
ties do  no  harm.  Purer  aluminium  sulphate  is  prepared  by  heating 


Aluminium.  349 

bauxite  with  soda  ash,  extracting  the  sodium  aluminate  formed  with 
water,  and  precipitating  the  aluminium,  as  the  hydroxide  with  carbon 
dioxide  gas.  The  relatively  pure  hydroxide  is  then  changed  into  sul- 
phate by  treatment  with  sulphuric  acid.  The  product,  known  as 
"concentrated  alum,1'  has  the  composition  expressed  by  the  formula 
A1.,(SO4)3  .  20  H2O,  though  separate  crystals  contain  only  eighteen 
molecules  of  water  of  crystallization.  By  boiling  cryolite  with  milk  of 
lime,  the  sodium  aluminate  thereby  formed  may  be  changed  into  "  con- 
centrated alum,"  as  described  above.  About  50,000  tons  of  "con- 
centrated alum  "  are  annually  produced  in  the  United  States. 

Alum.  — When  solutions  of  aluminium  sulphate  and  potas- 
sium sulphate  are  mixed  and  concentrated  by  evaporation, 
transparent,  colorless,  glassy  crystals  are  deposited.  This 
solid  is  potassium  alum,  or  simply  alum.  It  has  the  com- 
position represented  by  the  formula,  K2A12(SO4)4.  24  H2O, 
or  K2SO4,  A12(SO4)3 .  24  H2O,  and  is  sometimes  called  a 
double  salt.  It  is  the  type  of  a  class  of  similar  salts  called 
alums,  which  can  be  formed  by  crystallization  from  a 
mixture  of  aluminium  sulphate  and  an  alkaline  sulphate. 
Alums  are  very  soluble  in  water,  and  their  solutions  have 
an  acid  reaction  and  a  sweetish,  puckery  taste.  They 
crystallize  alike,  and  contain  twenty-four  molecules  of 
water  of  crystallization.  When  heated,  alums  lose  their 
water  of  crystallization  and  some  sulphuric  acid,  and  fall 
to  a  white  powder  or  porous  mass  known  as  burnt  alum. 
Potassium  alum  is  the  most  common,  but  ammonium  and 
sodium  alums  are  manufactured  and  used.  Sodium  alum 
is  an  ingredient  of  some  baking  powders.  Burnt  alum 
finds  application  as  a  medicine.  Alum  has  been  largely 
displaced  by  "  concentrated  alum,"  but  the  real  alum  is 
still  used  in  dyeing  and  printing  cloth,  in  tanning  and 
paper  making,  in  purifying  water  and  sewage,  as  a  medi- 
cine, for  hardening  plaster,  in  making  wood  and  cloth  fire- 
proof, and  in  preparing  other  aluminium  compounds. 


350  Descriptive  Chemistry. 

Alum  was  known  to  the  ancients,  who  used  it  in  dyeing  and  tanning, 
and  as  a  medicine.  It  was  first  manufactured  in  Europe,  about  the 
thirteenth  century,  from  native  alunite,  which  is  an  impure  sulphate  of 
aluminium,  potassium,  and  iron.  Alunite  and  alum  slates  or  shales  are 
now  used  to  some  extent,  but  most  of  the  alum  is  made  from  bauxite. 
Not  all  alums  contain  aluminium.  This  metal  may  be  replaced  by  iron, 
chromium,  manganese,  or  similar  metals,  producing  salhich  have 
the  same  general  properties  as  ordinary  alum. 
formula  of  alums  is  M2(SO4)3  .  X2SO4  .  24  H7O,  in 
aluminium,  iron,  chromium,  etc.,  and  X  a  metal  (or  group)  like  potas- 
sium, sodium,  ammonium.  Chrome  alum  (K.,Cr2(SO4)4 . 24  H.,0) 
belongs  to  this  class.  It  is  a  purple,  crystallized  solid.  The  other  alums 
have  a  limited,  industrial  application.  * 


Alums  and  other  aluminium  salts  are  used  as  mordants 
in  dyeing  and  calico  printing.  Some  dyes  must  be  fixed 
in  the  fabric  by  a  metallic  substance,  otherwise  the  color 
would  be  easily  removed.  The  cloth  to  be  dyed  or  printed 
is  impregnated  or  printed  with  the  mordant,  and  then 
heated  or  treated  with  some  substance  to  change  the  mor- 
dant into  an  insoluble  compound.  The  mordanted  cloth  is 
next  passed  through  a  vat  containing  the  solution  of  the 
dye,  which  unites  chemically  or  mechanically  (perhaps 
both)  with  the  metallic  compound,  forming  a  colored  com- 
pound. The  latter  is  called  a  "lake";  it  is  relatively  in- 
soluble, and  cannot  be  easily  washed  from  the  cloth,  i.e. 
it  is  a  fast  color.  Aluminium  acetate  or  "red  liquor"  and 
aluminium  sulphate,  besides  alum,  are  used  as  mordants 
for  cotton,  linen,  and  wool. 

Cryolite  is  a  white,  glassy,  crystallized  solid.  It  often 
resembles  clouded  ice,  and  its  name  means  "ice  stone." 
Its  composition  corresponds  to  the  formula  Na3AlF6  (or 
A1F3 .  3  NaF).  Small  fragments  melt  easily,  even  in  a 
candle  flame,  and  color  the  Bunsen  flame  yellow.  The 
only  locality  where  it  is  found  in  commercial  quantities  is 


Aluminium. 


351 


southern  Greenland,  which  yields  annually  about  10,000 
tons.  It  is  used  not  only  in  manufacturing  aluminium, 
but  as  a  source  of  alum  and  aluminium  hydroxide,  pure 
sodium  carbonate  and  hydroxide,  hydrofluoric  acid,  fluor- 
ides, and  one  kind  of  glass. 

Aluminium  Chloride  when  pure  is  a  white  powder,  but  it  is  often  a 
yellowish,  crystalline  mass  (A1C13  .  6  H2O).  It  is  prepared  by  heating 
powdered  aluminium  in  chlorine,  or  by  passing  chlorine  over  a  heated 
mixture  of  aluminium  oxide  and  carbon.  Exposed  to  the  air,  it  absorbs 
moisture  and  gives  off  fumes  of  hydrochloric  acid.  It  dissolves  in 
water  with  evolution  of  heat,  and  if  the  solution  is  heated,  hydrochloric 
acid  is  expelled,  owing  to  the  transformation  of  the  chloride  into  the 
hydroxide,  thus  — 

A1C13       +  3  H20  =        3  HC1        +        Al(OH), 
Aluminium       Water       Hydrochloric      Aluminium  Hy- 
Chloride  Acid  droxide 

This  salt  is  used  in  organic  chemistry. 

Clay  is  a  more  or  less  impure  aluminium  silicate,  formed 
by  the  slow  decomposition  of  rocks  containing  aluminium, 
especially  feldspar.  Pure  feldspar  is  a  silicate  of  alumin- 
ium and  sodium  or  potassium.  The  products  of  its  decom- 
position are  chiefly  an  insoluble  aluminium  silicate  and  a 
soluble  alkaline  silicate.  The  latter  is  washed  away.  The 
aluminium  silicate  which  remains  is  pure  clay  or  kaolin. 
The  latter  is  really  a  hydrous  silicate,  having  the  composi- 
tion corresponding  to  the  formula  Al2Si3O7,  2  H2O.  The 
composition  of  clay  varies,  because  it  is  seldom  formed 
from  pure  feldspar.  Most  kaolin  contains  particles  of  mica 
and  quartz.  Ordinary  clay  contains  many  impurities,  e.g. 
carbonates  of  calcium  and  magnesium,  quartz,  and  iron 
compounds.  Kaolin  is  a  white,  powdery  mass.  It  becomes 
slightly  plastic  when  wet,  and  can  therefore  be  molded 
into  various  shapes.  Ordinary  clay  is  very  plastic  when 


Descriptive  Chemistry. 

wet,  more  easily  fused  than  kaolin,  but  shrinks  consider- 
ably when  dried  and  burned ;  it  also  contains  iron  com- 
pounds, which  color  it  gray,  blue,  yellow,  brown,  and  red. 
All  clays  have  a  peculiar  clayey  odor  when  moist. 

Clay  is  the  basis  of  pottery,  of  which  there  are  three 
general  kinds  :  porcelain  or  china,  stoneware,  and  earthen- 
ware. 

Porcelain  is  the  finest  kind.  It  is  made  by  heating  to  a  high  tem- 
perature a  mixture  of  kaolin,  fine  sand,  and  some  fusible  substance,  such 
as  feldspar,  chalk,  or  gypsum.  The  mass  when  cool  is  hard,  dense,  white, 
and  translucent  (if  thin)  ;  it  is  not  easily  corroded  by  chemicals  (ex- 
cept fused  alkalies).  Although  it  is  not  very  porous,  its  surface  is 
glazed,  partly  for  protection,  partly  for  ornament.  This  is  done  by 
coating  it  with  a  mixture  similar  to  that  used  for  making  the  porcelain 
but  more  easily  fused,  and  then  heating  again  so  that  the  glaze  will 
penetrate  the  surface.  Stoneware  is  similar  to  porcelain,  but  coarser, 
because  the  materials  are  less  carefully  selected  and  prepared,  and  are 
not  heated  to  such  a  high  temperature.  The  best  grades  can  hardly  be 
distinguished  from  porcelain,  but  usually  stoneware  is  much  heavier 
and  thicker.  The  cheaper  kinds  are  made  into  jars,  jugs,  and  bottles, 
especially  large  ones  used  in  acid  manufactories.  Crockery  is  a  fine 
grade  of  stoneware,  though  the  best  crockery  is  much  like  porcelain. 
If  less  pure,  plastic  clay  is  used  and  heated  to  a  moderate  temperature, 
the  product  is  known  as  earthenware.  This  is  a  large  class  and  in- 
cludes majolica,  tiles,  terra  cotta,  jugs,  flowerpots,  clay  tobacco  pipes, 
drain  pipe,  and  bricks.  This  ware  is  porous  and  is  usually  glazed  by 
throwing  salt  into  the  baking  oven  just  before  the  operation  is  over. 
The  salt  volatilizes  arid  forms  a  fusible  sodium  aluminium  silicate  upon 
the  surface.  Cheap  bricks  are  made  from  very  impure  clay,  and  their 
red  color  is  due  to  iron  oxides  formed  from  the  iron  compounds  in  the 
unburned  clay.  Buff  bricks  are 'made  from  clay  containing  little  or  no 
iron,  and  clay  containing  silica  yields  fire-clay  bricks,  stove  linings, 
retorts,  and  crucibles. 

EXERCISES. 

1.  What  is  the  symbol  and  atomic  weight  of  aluminium  ? 

2.  Name  several  compounds  of  aluminium  and  discuss  their  occur- 
rence.    What  proportion  of  the  earth's  crust  is  aluminium  ? 


Aluminium.  353 

3.  State  briefly  the  history  of  aluminium. 

4.  Describe  the  metallurgy  of  aluminium  by  (#)  the  Hall  process, 
(£)  the  Heroult  process,  (V)  the  older  chemical  method. 

5.  Discuss  the  production  and  cost  of  aluminium. 

6.  (#)  Summarize  the  properties  of  aluminium.     (<£)  State  its  uses. 
(V)  Describe  its  alloys. 

7.  What  is  the  formula  and  chemical  name  of  alumina  ?     Describe 
its  preparation.     State  its  properties  and  uses.  » 

8.  State  the  properties  and  uses  of  corundum  and  emery.     Review 
carborundum  (see  Chapter  X). 

9.  Name  seven  gems  containing  aluminium.     Describe  them. 

10.  Describe   aluminium   hydroxide.      How   does   it   interact  with 
acids  and  with  alkalies  ? 

1 1 .  What  is  bauxite  ?     For  what  is  it  used  ? 

12.  Describe  aluminium  sulphate.      State  its  properties  and  uses. 
How  is  it  prepared  ?     What  is  "  alum  cake  "  ?     u  Alumino  ferric  cake  "  ? 
State  their  uses. 

13.  What  is  ordinary  alum  ?  How  is  it  manufactured  ?  State  the 
general  properties  and  uses  of  alums.  What  is  (a)  "concentrated 
alum,1'  and  (^)  burnt  alum  ? 

14.  Define  a  mordant.     Describe  its  use.     Name  several  mordants. 
What  is  (a)  a  "  lake,"  (b)  red  liquor  ? 

15.  What  is  the  general  formula  of  an  alum  ?      What  is  chrome 
alum  ? 

16.  Where  is  cryolite  found  ?    State  its  properties  and  uses.     What 
is  its  formula  ? 

17.  Describe  the  preparation  and  state  the  properties  of  aluminium 
chloride. 

18.  What  is  clay  ?    How  is  it  formed  ?  What  is  kaolin  ?    Describe 
(a)  ordinary  clay,  and  (6)  kaolin. 

19.  Describe  the  manufacture  of  (a)  porcelain,  (£)  stoneware,  and 
(V)    earthenware.      Give  an   example   of  each.     What   is  meant  by 
glazing  ? 

PROBLEMS. 

What  is  the  per  cent  of  aluminium  in  (a)  cryolite  (AlNa3F6), 
(£)  turquoise  (A1,P2O8  .  HrAl2O6  .  2  H2O),  (V)  corundum  (A12O3),  (W) 
aluminium  hydroxide  (A1(OH)3)  ? 


CHAPTER   XXVI. 
TIN  AND  LEAD. 

TIN  and  lead  are  familiar  metals.  They  have  similar 
and  useful  properties,  which  give  these  metals  and  their 
compounds  numerous  applications. 

TIN. 

Occurrence  of  Tin.  —  Metallic  tin  is  rarely  if  ever 
found.  Tin  dioxide  (cassiterite,  tin  stone,  SnO2)  is  the 
only  available  ore.  It  is  not  widely  distributed,  but  large 
deposits  are  found  in  England  (at  Cornwall),  Germany  (in 
Bohemia  and  Saxony),  Australia,  Tasmania,  and  the  East 
Indian  Islands,  especially  Banca  and  Billiton.  A  small 
quantity  is  found,  but  not  mined,  in  the  United  States. 

Tin  is  one  of  the  oldest  known  metals.  It  is  mentioned  in  the  Pen- 
tateuch, and  was  obtained  long  before  the  Christian  era  by  the  Phoeni- 
cians from  the  British  Isles,  which  were  called  Cassiterides  (from  the 
Greek  word  kassiteros,  meaning  tin).  Many  ancient  bronzes  contain 
tin.  The  alchemists  called  it  Jupiter,  and  used  the  metal  and  its  com- 
pounds. 

The  Latin  word  stannum  gives  us  the  symbol  Sn  and  the  terms 
stannous  and  stannic. 

Metallurgy  of  Tin.  —  If  the  tin  ore  contains  sulphur  or  arsenic,  these 
impurities  must  be  removed  by  roasting.  The  tin  oxide  is  then  reduced 
by  heating  it  with  coal  in  a  reverberatory  furnace  ;  the  simplest  equation 
for  this  change  is  — 

SnO2         +         C          =        Sn      +  CO2 

Tin  Dioxide         Carbon  Tin  Carbon  Dioxide 

354 


Tin  and  Lead.  355 

The  molten  tin  which  collects  at  the  bottom  of  the  furnace  is  drawn  off 
and  cast  into  bars  or  masses,  which  are  often  called  block  tin.  Usually 
it  is  purified  by  melting  it  slowly  on  a  hearth,  inclined  so  that  the  more 
easily  melted  tin  will  flow  down  the  hearth  and  leave  the  metallic  impuri- 
ties behind.  This  tin  may  be  further  purified  by  stirring  the  molten 
metal  with  a  wooden  pole,  or  by  holding  billets  of  wood  beneath  its  sur- 
face. The  impurities  which  are  oxidized  by  the  escaping  gases  collect 
as  a  scum  on  the  surface  and  are  removed. 

Properties  of  Tin.  —  Tin  is  a  white,  lustrous  metal, 
which  does  not  tarnish  easily  in  the  air.  It  is  soft  and 
malleable,  and  can  be  readily  cut  and  hammered.  It  is 
softer  than  zinc,  but  harder  than  lead.  Its  specific  gravity 
is  7.3.  Tin  may  be  obtained  in  the  crystalline  form,  and 
when  a  piece  of  such  tin  is  bent  it  makes  a  crackling 
sound,  which  is  caused  by  the  friction  of  these  crystals 
upon  one  another.  It  melts  at  about  232°  C,  and  when 
heated  to  a  higher  temperature'  it  burns,  forming  white  tin 
oxide  (SnO2).  The  physical  properties  of  tin,  like  those 
of  zinc,  vary  with  the  temperature.  Concentrated  hydro- 
chloric acid  changes  it  into  stannous  chloride  (SnCl2); 
treated  with  hot  concentrated  sulphuric  acid,  it  forms 
stannous  sulphate  (SnSO4)  and  sulphur  dioxide ;  and  com- 
mercial nitric  acid  oxidizes  it,  the  white,  solid  product 
being  known  as  metastannic  acid.  Zinc  precipitates  tin 
from  its  solutions  as  a  grayish  black,  spongy  mass,  which 
is  sometimes  filled  with  bright  scales. 

Uses  of  Tin. — Tin  is  so  permanent  in  air,  weak  acids 
(like  vinegar  and  fruit  acids),  and  alkalies  that  it  is  exten- 
sively used  as  a  protective  coating  for  metals.  Ordinary 
tinware  is  sheet  iron  coated  with  tin.  The  tin  plate 
(sheet  tin,  or  simply  "tin")  is  made  by  dipping  very  clean 
sheet  iron  into  molten  tin.  Tacks,  nails,  and  many  small 
iron  objects  are  similarly  tinned.  Copper  coated  with  tin 


356  Descriptive  Chemistry. 

is  made  into  vessels  for  cooking,  and  brass  coated  with 
tin  is  made  into  pins.  Large  quantities  of  tin  plate  are 
used  to  cover  roofs.  Tinned  iron  does  not  rust  until  the 
tin  is  worn  off  and  the  iron  exposed,  and  then  the  rusting 
proceeds  rapidly.  Tin  is  also  hammered  into  thin  sheets 
called  tin  foil,  though  much  of  the  tin  foil  now  used  con- 
tains lead.  Many  useful  alloys  contain  tin  as  an  essential 
ingredient.  During  the  last  few  years  the  annual  con- 
sumption of  tin  has  been  about  75,000  pounds. 

Alloys  of  tin  are  described  under  COPPER.  Those 
containing  a  minor  percentage  of  tin  are  .bronze,  gun 
metal,  bell  metal,  speculum  metal,  type  metal,  anti-friction 
metals,  and  fusible  alloys.  Britannia  metal  contains 
about  90  per  cent  tin,  8  per  cent  antimony,  and  the  rest 
mainly  copper.  It  is  a  white  metal,  and  was  formerly 
made  into  tableware.  White  metal  contains  less  tin  and 
more  antimony  than  Britannia,  though  the  composition 
varies.  It  resembles  Britannia.  The  harder  varieties  of 
'white  metal  are  used  as  parts  of  machinery,  and  the  softer 
kinds  are  made  into  ornaments  and  cheap  jewelry.  Pew- 
ter and  solder  contain  varying  proportions  of  tin  and  lead. 
Plumbers'  solder,  or  soft  solder,  is  about  one  third  tin  and 
two  thirds  lead.  It  is  harder  than  either  constituent,  but  it 
melts  at  a  lower  temperature.  Tin  amalgam  is  sometimes 
used  to  coat  mirrors. 

Compounds  of  Tin. — Tin  forms  two  series  of  compounds,  the  stan- 
nous  and  the  stannic.  Stannic  oxide  (SnO2)  has  already  been  men- 
tioned as  the  chief  ore  of  tin,  and  as  the  product  formed  when  tin  is 
burned.  The  artificial  oxide  is  faint  yellow  when  hot  and  white  when- 
cold.  The  native  oxide  is  a  brown  or  black,  lustrous,  and  often  crystal- 
lized solid.  Irregular  pebbles  called  stream  tin  occur  in  some  localities 
near  rivers.  Stannous  chloride  (SnCl.,)  is  formed  by  the  interaction 
of  hydrochloric  acid  and  tin.  From  the  concentrated  solution  a  green- 
ish salt  crystallizes  (SnCl2  .  i  H.,O),  known  as  tin  crystals  or  salt  of  tin. 


Tin  and  Lead.  357 

Stannous  chloride  passes  readily  into  stannic  chloride  (SnCl4)  when 
added  to  mercuric  chloride  solution.  The  simplest  equation  for  this 
change  is  — 

SnCl2         +        2  HgCl2         =         SnCl4         +         Hg2Cl2 
Stannous  Mercuric  Stannic  Mercurous 

Chloride  Chloride  Chloride  Chloride 

By  an  extension  of  the  simplest  idea  of  oxidation  and  reduction,  the 
stannous  chloride  in  the  change  is  said  to  be  oxidized  to  stannic  chlo- 
ride, but  it  reduced  the  mercuric  chloride  to  mercurous  chloride.  Stan- 
nous  chloride  is  often  used  as  a  reducing  agent  and  as  a  mordant  in 
dyeing  and  calico  printing.  Crystallized  stannic  chloride  (SnCl4 .  5  H2O), 
known  commercially  as  oxymuriate  of  tin,  is  also  used  as  a  mordant.  Tin 
mordants  produce  brilliant  colors.  Sodium  stannate  (Na2SnO3  .  3  H2O) 
is  extensively  used  to  prepare  cotton  cloth  for  printing. 

LEAD. 

Occurrence  of  Lead.  —  Metallic  lead  is  occasionally 
found  in  small  quantities.  The  most  abundant  ore  is  lead 
sulphide  (galena,  PbS).  Other  native  compounds,  formed 
by  the  alteration  of  galena,  are  the  carbonate  (cerussite, 
PbCO3),  the  sulphate  (anglesite,  PbSO4),  and  the  phos- 
phate (pyromorphite,  Pb5Cl(PO4)3).  Lead  compounds  are 
widely  distributed,  but  the  source  of  commercial  lead  is 
the  sulphide. 

Lead  has  been  used  by  civilized  people  since  the  dawn  of  history. 
The  Chinese  have  used  it  for  ages  to  line  chests  in  which  tea  is  stored 
and  transported.  The  Romans,  who  obtained  it  from  Spain,  called  it 
plumbum  nigrum,  i.e.  black  lead.  The  symbol  Pb  Qomes  from  plumbum. 
The  ancients  also  used  lead  compounds  (especially  the  carbonate  and 
red  oxide)  as  paints  and  cosmetics. 

•  The  annual  production  of  lead  has  increased  rapidly  during  the  last 
few  years,  and  in  1902  it  was  about  800,000  tons.  This  vast  amount 
comes  chiefly  from  the  United  States,  Spain,  Germany,  Mexico,  New 
South  Wales,  and  England.  The  United  States  in  1902  produced 
about  250,000  tons  of  lead  from  ores  found  mainly  in  the  Middle  West 
(Illinois,  Iowa,  Wisconsin,  and  Missouri),  Colorado.  Idaho,  and  Utah. 


358  Descriptive  Chemistry. 

Metallurgy  of  Lead.  —  Lead  is  readily  obtained  from  galena,  (i)  In 
the  reduction  process  the  ore  is  roasted  in  a  reverberatory  furnace  until 
a  part  of  the  sulphide  is  changed  into  lead  oxide  and  lead  sulphate. 
The  equations  for  these  changes  are  — 

2  PbS       4-       3  O2       =      2  PbO         +  2  SO2 

Lead  Sulphide         Oxygen         Lead  Oxide         Sulphur  Dioxide 

PbS  +         2O2  PbS04 

Lead  Sulphide  Oxygen  Lead  Sulphate 

The  air  is  then  shut  off  and  the  mixture  of  the  three  lead  compounds  is 
heated  to  a  higher  temperature.  By  this  operation  the  lead  sulphide 
interacts  with  the  other  lead  compounds,  forming  lead  and  sulphur  diox- 
ide, thus  — 

2  PbS         +       PbSO4        +     2  PbO     =  sPb  +         3  SO2 
Lead  Sulphide     Lead  Sulphate     Lead  Oxide     Lead      Sulphur  Dioxide 

(2)  Ores  poor  in  lead  are  sometimes  reduced  by  roasting  with  iron, 
which  combines  with  the  sulphur,  leaving  the  lead  free,  thus  — 

PbS  +       Fe       =     Pb         +  FeS 

Lead  Sulphide  Iron  Lead  Iron  Sulphide 

(3)  At  Niagara  Falls  lead  is  obtained  from  galena  by  electrolysis. 
Crushed  galena  is  made  the  cathode,  dilute  sulphuric  acid  is  the  electro- 
lyte, and  the  bottom  of  the  reduction  pan  is  the  anode.     The  sulphur 
is  changed  into  hydrogen  sulphide,  which  escapes  into  a  combustion 
chamber  where  its  sulphur  is  recovered  or  converted  into  sulphuric  acid. 
The  lead  remains  in  the  pan  as  a  spongy  mass.     The  silver,  which 
remains  in  the  lead  obtained  by  reduction,  is  extracted  by  the  Parkes 
process  (see  Silver). 

Properties  of  Lead.  —  Lead  is  a  bluish  metal.  When 
scraped  or  cut,  it  has  a  brilliant  luster,  which  soon  disap- 
pears, owing  to  the  formation  of  a  film  of  oxide.  This 
coating  protects  the  lead  from  further  change.  It  is  a  soft 
metal,  and  may  be  scratched  with  the  finger  nail.  It  dis- 
colors the  hands,  and  when  drawn  across  a  rough  surface 
it  leaves  a  black  mark.  For  this  reason  it  is  sometimes 


Tin  and  Lead.  359 

called  black  lead  (see  Graphite).  Lead  is  not  tough 
enough  to  be  readily  hammered  into  foil  or  drawn  into  fine 
wire,  but  it  can  be  rolled  into  sheets.  It  is  a  heavy  metal, 
its  specific  gravity  being  11.35;  with  the  exception  of 
mercury,  it  is  the  heaviest  of  the  familiar  metals.  It  melts 
at  326°  C,  or  about  100°  higher  than  tin  and  100°  lower 
than  zinc.  Lead,  when  heated  strongly  in  air,  changes 
into  an  oxide  (mainly  the  monoxide,  PbO).  Hydrochloric 
and  sulphuric  acids  have  little  effect  upon  compact  lead. 
Nitric  acid  changes  it  into  lead  nitrate  (Pb(NO3)2).  Acetic 
acid  (or  vinegar)  and  acids  from  fruits  and  vegetables 
change  it  into  soluble,  poisonous  compounds ;  hence  cheap 
tin-plated  vessels,  which  sometimes  contain  lead,  should 
never  be  used  in  cooking.  Zinc  and  iron  precipitate  lead 
from  its  solutions  as  a  grayish  mass,  which  often  has  a 
beautiful  treelike  appearance. 

Lead  in  Drinking  Water.  —  Lead  is  slowly  changed  into 
soluble  compounds  by  water  containing  carbon  dioxide, 
ammonia,  nitrates,  or  chlorides.  But  water  containing  sul- 
phates or  carbonates  forms  an  insoluble  coating  on  the 
lead,  thus  protecting  it  from  further  action.  All  lead  salts 
are  poisonous,  and  if  taken  into  the  system  they  will  slowly 
accumulate  and  ultimately  cause  serious  and  dangerous 
illness.  Water  suspected  of  attacking  lead  should  never 
be  drunk  after  it  has  been  standing  very  long  in  lead  pipes, 
but  should  be  allowed  to  flow  until  the  pipe  has  been  filled 
with  fresh  water.  Sometimes  the  water  cannot  be  drunk 
at  all.  The  city  of  Lowell,  Massachusetts,  recently  aban- 
doned one  source  of  its  water  supply  because  of  the  rapid 
solvent  action  of  the  water  upon  lead  pipes. 

Uses  of  Lead.  —  Lead  is  extensively  used  as  pipe,  be- 
cause it  can  be  made  into  indefinitely  long  pieces,  which 


Descriptive  Chemistry. 


can  be  easily  bent,  cut,  and  united  (by  solder).  The  pipe  is 
made  by  forcing  softened  lead  through  a  hole 
in  a  steel  plate  or  by  the  apparatus  shown 
in  Figure  69.  Lead  pipe  is  not  only  used 
to  convey  water  to  and  from  parts  of  build- 
ings, but  as  a  sheath  for  copper  wires,  both 
overhead  and  underground.  As  sheet  lead 
it  is  used  to  cover  roofs  and  to  line  sinks, 
cisterns,  and  the  cells  employed  in  many 
electrolytic  processes.  The  lead  chambers 
and  evaporating  pans  used  in  manufacturing 
sulphuric  acid  are  made  of  sheet  lead.  Shot 
and  bullets  are  lead  (alloyed  with  a  little 
arsenic).  Spongy  lead  is  used  in  preparing 

inthelongcylin-        ,  r 

der.cc,  is  forced    tne  plates  of  storage  batteries. 

K'tough  The  A11°ys  of  Lead  are  important.  Type 
metal  contains  70  to  80  per  cent  lead ;  the 


FIG.  69.—  Ap- 

Ing*  lead'  pipe! 
The  molten  lead 


the      space,     D, 

varied  insL  by  other  constituents  are  tin  and  antimony.  The 
the  steel  rod,  A.  latter  metal  expands  when  it  solidifies  and 
makes  the  face  of  the  type  sharp  and  clear. 
Solder,  pewter,  and  fusible  alloys  contain  lead  as  an 
essential  constituent  (see  Alloys  of  Tin).  Small  quantities 
are  found  in  brass  and  bronze. 

Lead  Oxides.  —  There  are  three  important  oxides.  Lead 
monoxide  (PbO)  is  a  yellowish  powder  known  as  massicot, 
or  a  buff-colored  crystalline  mass  called  litharge.  It  is 
formed  by  heating  lead  above  its  melting  point  in  a  cur- 
rent of  air.  It  is  made  this  way,  though  considerable  is 
obtained  as  a  by-product  in  separating  silver  from  lead 
(see  Cupellation).  Large  quantities  are  used  in  preparing 
some  oils  and  varnishes,  flint  glass,  other  lead  compounds, 
and  as  a  glaze.  Lead  tetroxide  (red  lead,  minium, 
Pb3O4)  is  a  red  powder,  'varying  somewhat  in  color  and 


Tin  and  Lead.  361 

composition.  It  is  prepared  by  heating  lead  (or  lead  mo- 
noxide) to  about  350°  C.  It  is  used  in  making  flint  glass. 
Pure  grades  are  made  into  artists'  paint,  but  the  cheap 
variety  is  used  to  paint  structural  iron  work  (bridges, 
gasometers,  etc.),  hulls  of  vessels,  and  agricultural  imple- 
ments. It  is  used  in  plumbing  and  gas  fitting  to  make 
joints  tight.  Orange  mineral  has  the  same  composition 
as  red  lead,  and  although  its  color  is  lighter,  its  uses  are  the 
same.  Lead  dioxide  (lead  peroxide,  PbO2),  is  a  brown 
powder  formed  by  treating  lead  tetroxide  with  nitric  acid. 
It  is  used  in  storage  batteries. 

Lead  Carbonate,  PbCO3,  is  found  native  as  the  trans- 
parent, crystallized  mineral  cerussite.  It  is  obtained  as  a 
white  powder  by  adding  ammonium  carbonate  solution  to 
lead  nitrate  solution.  Sodium  and  potassium  carbonates, 
however,  form  basic  lead  carbonates,  which  have  a  compo- 
sition depending  upon  the  temperature.  The  most  im- 
portant of  these  basic  carbonates  has  the  composition 
corresponding  to  the  formula  2  PbCO3 .  Pb(OH)2,  and  is 
known  as  white  lead.  It  is  a  heavy,  white  powder  which 
mixes  well  with  linseed  oil,  and  is  used  extensively  as  a 
white  paint  and  as  the  basis  of  many  colored  paints. 

White  lead  is  manufactured  by  several  processes.  The  Dutch  process 
is  the  oldest,  having  been  used  as  early  as  1622.  It  is  essentially  the 
same  to-day,  though  many  details  have  been  improved.  Perforated 
disks  of  lead  are  put  in  earthenware  pots  which  have  a  separate  com- 
partment at  the  bottom,  containing  a  weak  solution  of  acetic  acid 
(about  as  strong  as  vinegar).  These  pots  are  arranged  in  tiers  in 
a  large  brick  building,  and  spent  tan  bark  is  placed  between  each 
tier.  The  building  is  now  closed  except  openings  for  the  entrance  and 
exit  of  air  and  steam.  The  heat  volatilizes  the  acetic  acid  which  changes 
the  lead  into  a  lead  acetate.  The  tan  bark  ferments  and  liberates  car- 
bon dioxide,  which  changes  the  lead  acetate  into  basic  lead  carbonate 
or  white  lead.  The  whole  operation  requires  from  sixty  to  one  hun- 
dred days.  The  slowness  is  the  chief  objection  to  this  process.  In 


362 


Descriptive  Chemistry. 


the  German  process  acetic  acid  vapor,  steam,  and  carbon  dioxide  are 
forced  into  closed  chambers  in  which  sheets  of  lead  are  suspended.  It 
requires  about  five  weeks.  In  the  French  process  basic  lead  carbonate 
is  precipitated  from  a  basic  lead  acetate  by  carbon  dioxide.  Milner's 
process  is  a  modification  of  the  French  process.  Both  are  quicker  than 
the  Dutch  or  German  processes,  but  the  product  is  not  considered  so 
good.  An  electrolytic  process  has  recently  been  devised.  The  anode 
is  lead,  the  cathode  is  copper,  and  the  electrolyte  is  sodium  nitrate 
solution.  When  the  electric  current  is  passed,  (i)  nitric  acid  is  liber- 
ated at  the  anode,  and  changes  the  lead  into  lead  nitrate,  and  (2)  at 
the  cathode  sodium  is  formed,  which  decomposes  the  water,  thereby 
forming  sodium  hydroxide.  The  lead  nitrate  and  sodium  hydroxide 
solutions  interact,  forming  insoluble  lead  hydroxide  and  sodium  nitrate, 
thus  — 

Pb(NO3)2  +  2NaOH          =       Pb(OH)2       +       2  NaNO3 

Lead  Nitrate     Sodium  Hydroxide     Lead  Hydroxide     Sodium  Nitrate 

The  sodium  nitrate  is  left  in  the  cell  to  be  acted  upon  again,  but  the 
lead  hydroxide  is  changed  into  lead  carbonate  by  treatment  with  sodium 
bicarbonate.  This  process  is  rapid,  and  the  product  is  claimed  to  be 
as  good  as  white  lead  produced  by  other  processes.  White  lead  paint 
often  turns  dark  in  the  air,  owing  to  the  formation  of  lead  sulphide, 
which  is  black.  Its  extensive  use  is  largely  due  to  its  great  covering 
power,  i.e.  a  very  thin  layer  produces  a  perfectly  white  surface,  and 
therefore  less  paint  is  required  for  a  given  area.  It  is  often  adulterated 
with  zinc  oxide  and  barium  sulphate;  those  are  white  solids,  but  they 
are  cheaper  and  have  less  covering  power. 

Lead  Sulphide,  PbS.  —  Native  lead  sulphide  is  the  min- 
eral galena,  the  chief  ore  of  lead.     It  resembles  lead  in 


FIG.  70.  —  Galena  crystals  (cube,  octahedron  and  cube,  octahedron). 

appearance,   but  is  harder  and  is  usually  crystallized  as 
cubes,    octahedrons,  or  their  combinations  (Fig.  70).      It 


Tin  and  Lead.  363 

has  perfect  cubic  cleavage,  i.e.  it  breaks  into  cubes  or  frag- 
ments more  or  less  rectangular.  It  is  easily  changed  into 
lead  by  heating  it  alone  or  with  sodium  carbonate  on  char- 
coal. Lead  sulphide,  as  prepared  in  the  laboratory,  is  a 
black  solid. 

Black  lead  sulphide  is  readily  precipitated  from  a  lead  salt  solution 
by  hydrogen  sulphide.  Its  formation  is  the  test  for  lead.  It  is  changed 
into  lead  chloride  by  concentrated  hydrochloric  acid  and  into  lead  sul- 
phate by  concentrated  nitric  acid. 

Other  Compounds  of  Lead,  which  are  important,  are  the  chloride, 
sulphate,  nitrate,  chromate,  and  acetate.  Lead  chloride  (PbCl2)  is  a 
white  solid  formed  by  adding  hydrochloric  acid  or  a  soluble  chloride  to 
a  cold  solution  of  a  lead  salt.  It  dissolves  in  hot  water.  Lead  sul- 
phate (PbSO4)  is  a  white  solid,  formed  by  adding  sulphuric  acid  or  a 
soluble  sulphate  to  a  solution  of  a  lead  salt.  It  is  very  slightly  soluble 
in  water,  but  soluble  in  concentrated  sulphuric  acid,  hence  crude  sul- 
phuric acid  often  contains  lead  sulphate.  Lead  nitrate  (Pb(NO3)2)  is 
a  white  crystallized  solid  formed  by  dissolving  lead  (or  better,  lead  mo- 
noxide) in  nitric  acid.  When  heated,  it  decomposes  into  lead  oxide 
(PbO),  nitrogen  peroxide,  and  oxygen.  Lead  acetate  (Pb(C2H3O2)2) 
is  a  white,  crystallized  solid  formed  by  the  action  of  acetic  acid  upon 
lead  or  lead  oxide  (PbO) .  It  is  very  soluble  in  water  and  is  often 
called  "  sugar  of  lead.1' 

EXERCISES. 

1.  Name  the  chief  ore  of  tin.    Where  is  it  found?    What  is  "  stream 
tin"? 

2.  Give  briefly  the  history  of  tin.     What  is  its  symbol  ?    Why? 

3.  Describe  (a}  the  metallurgy  of  tin,  and  (6)  its  purification. 

4.  Summarize  the  properties  of  tin.     State  its*  uses. 

5.  What  is  "tin11?     Block  tin?     Tinfoil?     Tinware?     Sheet  tin? 
Tin  plate  ? 

6.  Describe  three  alloys  which  contain   large  proportions   of  tin. 
Name  several  alloys  containing  a  minor  proportion  of  tin. 

7.  Compare  native  and  artificial  tin  oxide  (SnO2). 

8.  What  is  the  formula  of  (a}  stannous  chloride,  and  (b)  stannic 
chloride?     What  is  their  chemical  relation?     State  the  use  of  each 
chloride.     What  other  names  has  stannous  chloride? 


364  Descriptive  Chemistry. 

9.   What  is  the  most  abundant  ore  of  lead?     Name  other  native 
compounds. 

10.  Give  a  brief  history  of  lead.     What  is  its  symbol?     Why? 

11.  Discuss  the  production  of  lead. 

12.  Describe  the  metallurgy  of  lead  by  (a)  the  reduction  process, 
(£)  roasting  with  iron,  (c)  electrolysis  of  galena. 

13.  Summarize  the  properties  of  lead. 

14.  State  the  uses  of  lead. 

15.  Discuss  the  relation  of  lead  to  water. 

1 6.  What  is  (a)  type  metal,  (6)  solder,  (c)  fusible  alloy? 

17.  Give  the  name  and  formula  of  the  oxides  of  lead. 

1 8.  Describe  the  preparation,  and  state  the  properties  and  uses  of 
(a)  litharge,  (#)  red  lead,  (c)  lead  peroxide. 

19.  What  is  white  lead?    Describe  its  preparation  by  (a)  the  Dutch 
method,  and  (£)  electrolysis  of  sodium  nitrate. 

20.  State  the  properties  and  uses  of  white  lead. 

21.  What  is  the  formula  and  chemical  name  of  galena?     Describe 
this  mineral.     Describe  the  corresponding  artificial  compound.     What 
is  the  test  for  lead? 

22.  Describe  the  following  salts  of  lead  :    (a)  chloride,  (b)  sulphate, 
(c)  nitrate,  (d)  acetate. 

PROBLEMS. 

1.  What  is  the  per  cent  of  lead  in  (a}  galena  (PbS),  (£)  cerussite 
(PbCO8),  (c)  anglesite  (PbSO4),  (d)  lead  acetate  (Pb(C2H3O2)2 .  3  H2O)  ? 

2.  How  much  litharge  may  be  made  from  40.5  gm.  of  lead?     (As- 
sume Pb  +  O  =  PbO.) 

3.  What  is  the  per  cent  of  tin  in  (a)  tinstone  (SnO2),  (b)  stannous 
chloride  (SnCl2),  (c)  stannic  chloride  (SnCl4)? 


CHAPTER    XXVII. 
CHROMIUM  AND  MANGANESE. 

THESE  elements  do  not  belong  to  the  same  group,  but 
they  have  several  common  properties  and  form  analogous 
compounds. 

CHROMIUM. 

Occurrence  of  Chromium.  —  Metallic  chromium  is  never 
found  free.  Its  chief  ore  is  an  oxide  (chromite,  chrome 
iron  ore,  FeCr2O4).  Native  lead  chromate  (crocoite  or 
crocoisite,  PbCrO4)  is  less  common.  Traces  of  chromium 
occur  in  many  green  minerals  and  rocks,  e.g.  emerald  and 
serpentine,  and  verde  antique  marble. 

Chromite  is  mined  chiefly  in  Greece,  New  Caledonia,  New  South 
Wales,  Turkey,  and  Canada.  The  total  annual  production  is  about 
30,000  tons. 

The  word  chromium  comes  from  the  Greek  word  chroma,  meaning 
color,  and  emphasizes  the  fact  that  most  chromium  compounds  have 
decided  colors. 

Preparation,  Properties,  and  Uses.  —  Chromium  was  a  rare  metal 
until  Moissan  prepared  it,  in  1894,  in  the  electric  furnace.  Now  it  is 
produced  in  quantities  by  heating  a  mixture  of  chromite  and  carbon  in 
an  electric  furnace.  The  crude  chromium  is  refined  by  fusing  it  with 
lime.  Very  pure  chromium  is  also  prepared  by  reducing  chromic  oxide 
with  aluminium  powder. 

Chromium  is  a  lustrous  gray  metal.  It  takes  a  good  polish,  which  is 
not  removed  by  exposure  to  air.  It  is  hard,  but  it  can  be  filed  and  pol- 
ished without  difficulty.  Its  specific  gravity  is  about  6.9.  It  is  not 
attracted  by  a  magnet.  It  can  be  fused  only  in  the  electric  furnace. 

Chromium  is  used  to  harden  the  steel,  which  is  to  be  made  into 
armor,  projectiles,  safes,  and  vaults,  and  parts  of  machines  used  to 

365 


366  Descriptive  Chemistry. 

crush  gold-bearing  quartz.  This  hardened  steel  is  called  chrome  steel. 
The  commercial  form  of  chromium  is  an  alloy  of  65  to  80  per  cert 
chromium,  a  little  carbon,  and  the  rest  iron  ;  this  alloy  is  called  ferro- 
chrome. 

Compounds  of  Chromium  are  numerous,  some  are  com- 
plex, many  pass  readily  into  one  another,  and  a  few  have 
industrial  applications.  The  most  important  are  potassium 
chromate,  potassium  dichromate,  chrome  alum,  and  lead 
chromate. 

Potassium  Chromate  (K2CrO4)  and  Potassium  Dichro- 
mate (or  Bichromate,  K2Cr2O7).  —  These  compounds  are 
manufactured  from  chrome  iron  ore.  The  crushed  ore  is 
mixed  with  lime  and  potassium  carbonate,  and  roasted  in 
a  reverberatory  furnace ;  air  is  freely  admitted  and  the 
mass  is  frequently  raked.  By  this  operation  the  ore  is 
oxidized  into  a  mixture  of  calcium  and  potassium  chro- 
mates.  The  mass  is  cooled,  pulverized,  and  treated  with 
a  hot  solution  of  potassium  sulphate,  which  changes  the 
calcium  chromate  into  potassium  chromate.  The  clear, 
saturated  solution  of  potassium  chromate  is  changed  by 
sulphuric  acid  into  potassium  dichromate ;  the  latter  is 
purified  by  recrystallization  from  water.  Potassium  chro- 
mate is  a  lemon-yellow,  crystallized  solid,  very  soluble  in 
water.  Acids  change  it  into  the  dichromate,  thus  — 

2  K2CrO4  +  H2SO4  =  K2Cr2O7  +  K2SO4  +  H2O 
Potassium  Sulphuric  Potassium  Potassium  Water 
Chromate  Acid  Dichromate  Sulphate 

Potassium  Dichromate  is  a  red  solid  which  forms  large 
crystals.  It  is  less  soluble  in  water  than  potassium  chro- 
mate. Alkalies  change  it  into  a  chromate,  thus  — 

K2Cr2O7    +    2KOH   =   2  K2CrO4  +   H2O 

Potassium  Potassium  Potassium  Water 

Dichromate         Hydroxide  Chromate 


Chromium  and   Manganese.  367 

Potassium  dichromate  is  used  in  dyeing,  calico  printing, 
and  tanning,  in  bleaching  oils,  and  in  manufacturing  other 
chromium  compounds  and  dyestuffs.  Its  uses  depend 
mainly  upon  the  fact  that  it  is  an  oxidizing  agent.  When 
hydrochloric  acid  is  added  to  potassium  dichromate,  oxy-r 
gen  from  the  dichromate  withdraws  hydrogen  from  the 
acid  and  liberates  free  chlorine,  thus  — 

K2Cr2O7  -f  14  HC1  =  2  KC1  +  2  CrCl3  +  3  C12  +  7  H2O 

Potassium  Di-      Hydrochloric    Potassium        Chromic        Chlorine  Water 

chromate  Acid  Chloride         Chloride 

If  an  oxidizable  substance  is  present,  such  as  organic  mat- 
ter, alcohol,  or  a  ferrous  compound,  it  is  quickly  oxidized. 

Potassium  chromate  is  also  formed  as  a  yellow  mass  by  fusing  on 
porcelain  or  platinum  a  mixture  of  a  chromium  compound,  potassium 
carbonate,  and  potassium  nitrate.  When  the  mass  is  boiled  with  acetic 
acid  to  decompose  the  carbonate  and  expel  carbon  dioxide,  and  then 
added  to  a  lead  salt  solution,  yellow  lead  chromate  is  formed.  This 
experiment  is  often  used  as  a  test  for  chromium. 

Chrome  Alum,  K2Cr2  (SO4)4 .  24  H2O,  is  a  purple,  crys- 
tallized solid.  It  is  analogous  in  composition  and  similar 
in  properties  to  ordinary  alum,  but  it  contains  chromium 
instead  of  aluminium.  It  can  be  prepared  by  mixing 
potassium  and  chromium  sulphates  in  the  proper  propor- 
tion, or  by  passing  sulphur  dioxide  into  a  solution  of 
potassium  dichromate  containing  sulphuric  acid.  The 
commercial  substance  is  a  by-product  obtained  in  the 
manufacture  of  alizarine,  a  dye  which  yields  magnificent 
colors.  Chrome  alum  is  used  as  a  mordant  in  dyeing  and 
calico  printing,  and  in  tanning. 

Lead  Chromate,  PbCrO4,  is  a  bright  yellow  solid,  formed 
by  adding  potassium  chromate  or  dichromate  to  a  solution 
of  lead  salt:  It  is  known  as  chrome  yellow  and  is  used 
as  the  basis  of  yellow  paint  When  boiled  with  sodium 


370  Descriptive  Chemistry. 

called  black  oxide  of  manganese.  When  heated  it  yields 
oxygen ;  and  when  heated  with  hydrochloric  acid  the  two 
compounds  interact,  forming  manganous  chloride,  chlorine, 
and  water,  thus  — 

MnO2     +     4HC1     =     MnCl2     +     Cla     +   H2O 
Manganese       Hydrochloric       Manganese        Chlorine         Water 
Dioxide  Acid  Chloride 

It  colors  glass  and  borax  a  beautiful  amethyst,  and"  is  often 
added  to  common  glass  to  neutralize  the  green  color. 
Enormous  quantities  are  used  in  the  manufacture  of  oxy- 
gen, chlorine,  glass,  and  manganese  alloys  and  compounds. 

The  manganese  dioxide  used  in  the  manufacture  of  chlorine  is  recov- 
ered by  the  Weldon  process.  The  impure  manganous  chloride  solu- 
tion from  the  chlorine  still  is  treated  with  calcium  carbonate  to  neutralize 
free  acid  and  precipitate  any  iron  present.  Lime  is  added  to  the  clear 
solution  of  manganous  chloride,  and  air  is  blown  into  the  mixture.  The 
manganous  chloride  is  changed  into  manganous  hydroxide  (Mn(OH).,), 
which  interacts  with  the  oxygen  (of  the  air)  and  lime,  forming  chiefly 
calcium  manganite  (CaMnO3,  or  CaO  .  MnO2).  After  this  mixture  has 
settled,  the  calcium  chloride  is  drawn  off,  and  the  manganese  compound, 
which  is  called  "  Weldon  mud,"  is  used  to  generate  more  chlorine. 

Manganese  dioxide  was  used  by  the  ancients  to  decolorize  glass,  but 
its  nature  was  misunderstood.  They  confused  it  with  an  iron  oxide 
called  magnesia  stone,  and  the  alchemists  in  the  Middle  Ages  gave 
the  name  magnesia  to  this  manganese  dioxide.  Later  they  called  it 
magnesia  nigra,  or  black  magnesia,  to  distinguish  it  from  magnesia  alba, 
or  white  magnesia  (MgO),  supposing  that  the  two  were  related.  Man- 
ganese was  isolated  in  1774,  and  later  was  given  the  specific  name 
manganesium,  which  was  soon  shortened  to  manganese. 

Potassium  Permanganate,  KMnO4,  is  a  dark  purple, 
glistening,  crystallized  solid,  though  the  crystals  sometimes 
appear  black,  with  a  greenish  luster.  It  is  very  soluble  in 
water,  and  the  solution  is  red,  purple,  or  black,  according 
to  the  concentration.  Potassium  permanganate  gives  up 
its  oxygen  readily  and  is  used  as  an  oxidizing  agent  in  the 


Chromium  and  Manganese.  371 

laboratory  and  on  a  large  scale  to  purify  stagnant  water 
and  sewage.  It  is  such  a  powerful  oxidizing  agent  that  it 
cannot  be  filtered  through  paper,  but  only  through  asbestos 
or  spun  glass.  It  is  also  used  as  a  disinfectant,  as  a  medi- 
cine, in  bleaching  and  dyeing,  in  coloring  wood  brown,  and 
in  purifying  gases,  such  as  hydrogen,  ammonia,  and  carbon 
dioxide. 

Potassium  permanganate  is  manufactured  by  oxidizing  a  mixture  of 
manganese  dioxide  and  potassium  hydroxide,  and  treating  the  resulting 
potassium  manganate  with  sulphuric  acid,  carbon  dioxide,  or  chlorine. 
The  essential  reactions  are  represented  thus  — 

MnO,       +       2KOH     +  O  =    K,MnO4     +  H2O 
Manganese        Potassium  Potassium 

Dioxide  Hydroxide  Manganate 

3  K2MnO4  +  2  CO2  =  2  KMnO4  +  K2CO3  +  MnO2 

Potassium  Permanganate 

The  uses  of  potassium  permanganate  depend  mainly  upon  its  oxidiz- 
ing power.  With  sulphuric  acid  the  action  is  represented  thus  — 

2KMnO4    +   3H2SO4  =    50    +    2  MnSO4  +   K2SO4  +   3  H2O 
Potassium        Sulphuric     Oxygen    Manganese    Potassium     Water 
Permanganate        Acid  Sulphate        Sulphate 

The  liberated  oxygen  attacks  at  once  any  organic  matter  present,  and 
the  solution  becomes  brown  or  colorless,  owing  to  the  decomposition 
of  the  potassium  permanganate  into  colorless  compounds. 

Compounds  of  Manganese,  like  those  of  chromium,  are  numerous, 
often  complex,  and  closely  related.  There  are  four  oxides  besides 
manganese  dioxide.  Three  manganous  compounds  are  important,  the 
chloride  (MnCL,),  the  sulphate  (MnSO4),  and  the  sulphide  (MnS). 
The  chloride  and  sulphate  are  pink,  crystallized  salts,  and  the  sulphide 
is  a  flesh-colored  precipitate  formed  by  adding  ammonium  sulphide  to 
the  solution  of  a  manganous  salt,  thus  distinguishing  it  from  all  other 
sulphides.  Manganates  are  salts  of  the  hypothetical  manganic  acid 
(H2MnO4).  They  are  analogous  to  chromates,  and  the  manganese  in 
them  acts  as  a  non-metal.  Potassium  manganate  is  obtained  as  a 
green  mass  by  fusing  a  mixture  of  a  manganese  compound,  potassium 


372  Descriptive  Chemistry. 

hydroxide  (or  carbonate),  and  potassium  nitrate.  Its  formation  on  a 
small  scale  constitutes  the  test  for  manganese.  Sodium  manganate 
is  used  in  solution  as  a  disinfectant. 


EXERCISES. 

i  .   What  is  the  symbol  of  chromium  and  of  manganese  ?   Why  is  each 
element  so  named? 

2.  What  is  the  chief  ore  of  chromium?    Where  is  it  found?    What 
other  minerals  contain  chromium? 

3.  Describe  the  preparation  of  chromium.     State  its  properties  and 
uses.     What  is  chrome  steel?     Ferrochrome? 

4.  Describe  the  manufacture  of  (a)  potassium  chromate,  and  (£)  po- 
tassium dichromate.     State  their  properties  and  uses.     What   is   the 
formula  of  each? 

5.  What  are  the  tests  for  chromium? 

6.  Describe  chrome  alum.     How  is  it  made?     State  its  uses.     How 
does  it  differ  from  ordinary  alum? 

7.  Describe  lead  chromate.    How  is  it  formed  ?    For  what  is  it  used  ? 

8.  In  what  two  ways  does  chromium  act  in  its  compounds?     What 
is  chromic  oxide?     For  what  is  it  used?     What  is  chromium  trioxide? 
How  is  it  related  to  potassium  dichromate? 

9.  Name  several  ores  of  manganese.     What  is  the  chief  ore  ?     Dis- 
cuss the  production  of  manganese  ores. 

10.  Describe  the  preparation,  and  state  the  properties  of  manganese. 

11.  What  is  spiegel  iron?     Ferromanganese?     State  their  uses. 

12.  Describe  manganese  dioxide.     State   its   properties   and   uses. 
How  is  it  recovered  by  the  Weldon  process?     What  is  the  common 
name  of  manganese  dioxide?     Why  is  it  so  called? 

PROBLEMS. 

1.  What  is  the  per  cent  of  chromium  in  (a)  lead  chromate  (PbCrO4), 
(£)  chrome  ironstone  (Cr2O3  .  FeO),  (c)  chromic  oxide  (Cr2O3)  ? 

2.  What  is  the  per  cent  of  manganese  in  (<z)  manganese  dioxide 
(MnO9),    (£)    manganese     sulphide    (MnS),     (<:)    manganese    alum 
K2Mn2(S04)4.24H20)? 

3.  How  much  manganese  ore  containing  85  per  cent  of  manganese 
dioxide  is  needed  to  prepare  300  Ib.  of  chlorine?     (Assume  MnO2  + 
4HC1  =  C12  +  MnCl,  +  2  H2O.) 


CHAPTER  XXVIII. 
IRON,  NICKEL,  AND  COBALT. 

Introduction.  —  These  three  elements  form  a  natural 
group.  Their  properties  are  similar.  Cobalt  and  nickel 
are  very  closely  related  and  are  seldom  found  alone.  Iron 
resembles  manganese  and  chromium. 

IRON. 

Iron  is  the  most  useful  of  all  metals.  It  has  been  known 
for  ages,  and  has  been  indispensable  in  the  development  of 
the  human  race. 

The  symbol  of  iron,  Fe,  is  from  the  Latin  wordferrum.  Yromferrum 
are  derived  the  forms  ferri-  and  ferro-  (found  in  such  words  as  ferricya- 
nide,  ferro  manganese,  ferrocyanide,  etc.),  and  the  terms  ferrous  and 
ferric. 

Occurrence  of  Iron.  —  Uncombined  iron  is  found  only 
in  meteorites,  which  fall  upon  the  earth  from  remote 
regions  in  space,  and  in  a  very  few  rocks.  Combined  iron 
is  abundant  and  widely  distributed.  It  is  found  in  most 
rocks  and  many  minerals,  in  the  soil,  in  springs  and  nat- 
ural waters,  in  chlorophyll  —  the  green  coloring  matter  of 
plants,  —  and  in  haemoglobin  —  the  red  coloring  matter 
of  blood.  The  chief  ores  of  iron  are  hematite  (Fe2O3), 
limonite  (Fe2O3.  Fe2(OH))6,  magnetite  (Fe3O4),  and  sider- 
ite  (FeC03). 

Other  abundant  compounds  of  iron  not  used  as  a  source  of  the  metal 
are  pyrites  (FeS2),  pyrrhotite  (varying  from  Fe6S7  to  FenS12),  and  the 
copper-iron  sulphides  (chalcopyrite,  CuFeS2,  and  bornite,  Cu3FeS3). 

373 


372  Descriptive  Chemistry. 

hydroxide  (or  carbonate),  and  potassium  nitrate.  Its  formation  on  a 
small  scale  constitutes  the  test  for  manganese.  Sodium  manganate 
(NaMnOJ  is  used  in  solution  as  a  disinfectant. 

EXERCISES. 

1 .  What  is  the  symbol  of  chromium  and  of  manganese  ?   Why  is  each 
element  so  named  ? 

2.  What  is  the  chief  ore  of  chromium?    Where  is  it  found?    What 
other  minerals  contain  chromium? 

3.  Describe  the  preparation  of  chromium.     State  its  properties  and 
uses.     What  is  chrome  steel?     Ferrochrome? 

4.  Describe  the  manufacture  of  (#)  potassium  chromate,  and  (£)  po- 
tassium dichromate.     State  their  properties  and  uses.     What   is   the 
formula  of  each? 

5.  What  are  the  tests  for  chromium? 

6.  Describe  chrome  alum.     How  is  it  made?     State  its  uses.     How 
does  it  differ  from  ordinary  alum  ? 

7.  Describe  lead  chromate.    How  is  it  formed  ?    For  what  is  it  used  ? 

8.  In  what  two  ways  does  chromium  act  in  its  compounds?     What 
is  chromic  oxide?     For  what  is  it  used?     What  is  chromium  trioxide? 
How  is  it  related  to  potassium  dichromate? 

9.  Name  several  ores  of  manganese.     What  is  the  chief  ore  ?     Dis- 
cuss the  production  of  manganese  ores. 

10.  Describe  the  preparation,  and  state  the  properties  of  manganese. 

11.  What  is  spiegel  iron?     Ferromanganese ?     State  their  uses. 

12.  Describe  manganese  dioxide.     State   its   properties   and   uses. 
How  is  it  recovered  by  the  Weldon  process?     What  is  the  common 
name  of  manganese  dioxide?     Why  is  it  so  called? 

PROBLEMS. 

1.  What  is  the  per  cent  of  chromium  in  (a}  lead  chromate  (PbCrO4), 
(b)  chrome  ironstone  (Cr2O3  .  FeO),  (<r)  chromic  oxide  (Cr2O3)  ? 

2.  What  is  the  per  cent  of  manganese  in  (a}  manganese  dioxide 
(MnO2),     (<£)    manganese     sulphide    (MnS),     (c)    manganese     alum 
K2Mn2(SO4)4.24H,0)? 

3.  How  much  manganese  ore  containing  85  per  cent  of  manganese 
dioxide  is  needed  to  prepare  300  Ib.  of  chlorine?     (Assume  MnO2  + 
4HC1  =  C12  +  MnCl,  +  2  H2O.) 


CHAPTER  XXVIII. 
IRON,  NICKEL,  AND  COBALT. 

Introduction.  —  These  three  elements  form  a  natural 
group.  Their  properties  are  similar.  Cobalt  and  nickel 
are  very  closely  related  and  are  seldom  found  alone.  Iron 
resembles  manganese  and  chromium. 

IRON. 

Iron  is  the  most  useful  of  all  metals.  It  has  been  known 
for  ages,  and  has  been  indispensable  in  the  development  of 
the  human  race. 

The  symbol  of  iron,  Fe,  is  from  the  Latin  word  fer 'rum.  Yromferrum 
are  derived  the  forms  ferri-  and  ferro-  (found  in  such  words  as  ferricya- 
nide,  ferro  manganese,  ferrocyanide,  etc.),  and  the  terms  ferrous  and 
ferric. 

Occurrence  of  Iron.  —  Uncombined  iron  is  found  only 
in  meteorites,  which  fall  upon  the  earth  from  remote 
regions  in  space,  and  in  a  very  few  rocks.  Combined  iron 
is  abundant  and  widely  distributed.  It  is  found  in  most 
rocks  and  many  minerals,  in  the  soil,  in  springs  and  nat- 
ural waters,  in  chlorophyll  —  the  green  coloring  matter  of 
plants,  —  and  in  haemoglobin  —  the  red  coloring  matter 
of  blood.  The  chief  ores  of  iron  are  hematite  (Fe2O3), 
limonite  (Fe2O3  .  Fe2(OH))6,  magnetite  (Fe3O4),  and  sider- 
ite  (FeC03). 

Other  abundant  compounds  of  iron  not  used  as  a  source  of  the  metal 
are  pyrites  (FeS2),  pyrrhotite  (varying  from  FefiS7  to  FenS12).  and  the 
copper-iron  sulphides  (chalcopyrite,  CuFeS2,  and  bornite,  Cu3FeS3). 

373 


374 


Descriptive  Chemistry. 


The  United  States  leads  the  world  in  the  production  of  iron  ore,  the 
annual  output  for  the  last  few  years  being  over  25,000,000  tons.  This 
vast  quantity  comes  from  twenty-five  different  states,  but  the  bulk  is 
mined  in  Minnesota,  Michigan,  Alabama,  Wisconsin,  Tennessee,  Vir- 
ginia and  West  Virginia,  and  Colorado.  The  most  abundant  ore  is  the 
red  hematite,  which  comes  chiefly  from  the  Lake  Superior  region  (Fig. 
71)  ;  large  quantities  are  mined  in  Alabama  and  Tennessee.  The 
latter  states,  together  with  Virginia  and  West  Virginia,  furnish  most  of 
the  limonite  or  brown  iron  ore.  Pennsylvania,  New  Jersey,  and  New 
York  contribute  most  of  the  magnetite,  though  some  is  mined  also  in 


FIG.  71.  —  Deposits  of  iron  and  copper  near  Lake  Superior.  No.  4  is  the  cop- 
per region.  The  iron  regions,  known  as  ranges,  are  Marquette  (i),  Menominee 
(2),  Gogebic  (3),  Vermilion  (5),  Mesabi  (6). 

Michigan.  The  carbonate  ores,  which  constitute  less  than  one  per  cent 
of  the  output,  come  mainly  from  Ohio,  Maryland,  and  New  York.  Im- 
provements in  the  machinery  and  methods  used  in  mining  and  trans- 
porting iron  ore  have  reduced  its  cost  and  facilitated  its  production. 
Thus,  at  an  incredibly  small  expense,  ore  from  the  Lake  Superior  region 
is  raised  from  open  pits  by  steam  shovels,  dumped  into  large  cars,  car- 
ried to  shipping  ports  on  the  lakes,  dumped  again  into  huge  bunkers, 
dropped  down  chutes  into  big  freight  steamers  (many  of  which  hold 
6000  tons),  which  carry  it  to  South  Chicago  and  Milwaukee,  though 
over  two  thirds  is  received  at  ports  on  the  south  shore  of  Lake  Erie 


Iron,  Nickel,  and  Cobalt. 


375 


and  forwarded  by  rail  to  Pittsburg,  Pennsylvania.  This  city  is  the  great 
center  of  the  iron  and  steel  industries.  Birmingham,  Alabama,  is  the 
center  of  the  industry  in  the  South,  because  near  it  the  necessary  ore, 
coal,  and  limestone  are  con- 
veniently located. 


Metallurgy  of  Iron. 

—  Iron  is  extracted  most 
easily  from  its  oxides. 
The  ores,  whatever  their 
character,  are  first 
crushed  and  roasted  to 
change  them  into  ferric 
oxide  (Fe2O3)  as  far  as 
possible,  and  to  make 
the  raw  material  porous. 
Thus  prepared,  the  ore 
is  smelted  with  coke  (or 
coal)  and  limestone  in 
a  blast  furnace.  The 
carbon  reduces  the  oxide 
to  metallic  iron,  which 
collects  as  a  liquid  at  the 
bottom  of  the  furnace 
beneath  the  slag  formed 
by  the  limestone  and 
impurities.  The  blast 
furnace  (Fig.  72)  is  a 
huge  tower,  from  forty 
to  ninety  feet  high  and 
from  fourteen  to  seven- 
teen feet  in  diameter  at 
the  largest  part;  but  it 
is  narrower  at  the  top 


FIG.  72.  —  Blast  furnace.  A,  throat;  B, 
bosh ;  C,  crucible  where  the  melted  iron  col- 
lects ;  D,  pipes  for  hot  air  blast ;  E,  escape 
pipe  for  gases  which  do  not  escape  through 
the  "  down  comer  " ;  G,  cup ;  //,  cone ;  N, 
trough  for  drawing  off  slag ;  T,  tuyere ;  /,  hole 
through  which  iron  is  withdrawn. 

and  bottom  than  in  the  middle. 


376  Descriptive  Chemistry. 

It  is  built  of  masonry  and  iron,  and  lined  with  fire  brick. 
Pipes  at  the  bottom,  called  tuyeres,  allow  large  quantities 
of  hot  air  to  be  forced  into  the  furnace  and  up  through  the 
contents,  thereby  producing  the  high  temperature  required 
in  the  melting ;  while  another  pipe  at  the  top  not  only  per- 
mits the  escape  of  hot  gaseous  products,  but  conducts 
them  into  a  series  of  pipes  which  lead  to  different  parts  of 
the  plant,  where  the  hot  gases  are  utilized  as  fuel.  The 
blast  pipes  correspond  to  the  bellows  used  by  a  blacksmith, 
and  the  exit  pipe  to  a  chimney,  except  that  gases  escaping 
through  chimneys  are  usually  wasted. 

When  the  furnace  has  been  heated  to  the  proper  tem- 
perature, or  is  already  in  operation,  the  ore,  coke,  lime- 
stone, etc.,  are  carried  to  the  top  of  the  furnace  by 
machinery  and  introduced  into  the  furnace  by  dumping 
them  upon  the  cone-shaped  cover;  their  weight  lowers 
the  cover,  which  flies  back  tightly  into  place  after  the 
materials  roll  into  the  furnace.  The  charge  consists  of 
alternate  layers  of  ore,  fuel,  and  flux.  The  fuel  is  coke,  or 
coke  mixed  with  coal.  The  flux  varies  with  the  ore,  but 
it  is  usually  limestone,  though  feldspar  and  sand  are  used 
if  the  ore  contains  lime  compounds.  The  object  of  the 
flux  is  twofold,  (i)  It  removes  the  impurities  from  the 
charge  in  the  form  of  a  fusible  glass  called  slag  or  cinder, 
and  (2)  thereby  prevents  the  reduced  iron  from  reuniting 
with  oxygen  of  the  air  which  is  being  constantly  blown  in. 
As  the  smelting  proceeds,  the  iron  falls  through  the  slag 
to  the  bottom  of  the  furnace,  where  both  are  drawn  off 
through  separate  openings.  Fresh  charges  of  definite 
weight  are  added  at  regular  intervals,  and  the  whole  oper- 
ation continues  without  interruption  for  months  or  even 
years. 

The  iron  from  the  furnace  is  usually  poured  into  molds 


Iron,  Nickel,  and  Cobalt.  377 

of  sand  and  allowed  to  solidify.  Such  iron  is  called  pig 
iron  or  cast  iron.  A  single  large  furnace  will  produce  in 
a  day  about  500  tons  of  pig  iron.  In  some  plants  the 
molten  iron  is  run  into  huge  vessels,  called  converters,  and 
made  directly  into  steel  (see  below). 

The  chemical  changes  involved  in  the  metallurgy  of  iron  are  numer- 
ous and  complicated.  In  general,  the  iron  oxide  is  reduced  to  metallic 
iron  largely  by  carbon  monoxide.  The  carbon  of  the  fuel  at  first  forms 
carbon  dioxide  with  the  oxygen  of  the  air  blast.  But  the  dioxide  is 
soon  reduced  by  the  hot  carbon  to  the  monoxide,  which  interacts  with 
the  ore,  thus  — 

Fe20s     +     3  CO     =     2Fe     +     3  CO2 
Ferric  Carbon  Iron  Carbon 

Oxide        Monoxide  Dioxide 

Considerable  carbon  monoxide  escapes,  however,  in  the  waste  gas.  At 
this  stage  the  iron  becomes  porous,  and  is  doubtless  prevented  from  re- 
oxidation  by  the  carbon  dioxide  liberated  from  the  decomposed  lime- 
stone. As  the  spongy  iron  sinks  into  the  hotter  part  of  the  furnace,  it 
combines  with  carbon  to  some  extent,  finally  melts,  sinks  through  the 
slag,  and  accumulates  at  the  bottom,  or  hearth,  of  the  furnace.  The 
iron  obtained  in  this  way  contains  small  amounts  of  carbon,  sulphur, 
phosphorus,  silicon,  and  manganese. 

About  300  blast  furnaces  were  in  operation  in  the  United  States  in 
1902,  and  the  number  is  increasing.  They  consumed  over  30,000,000 
tons  of  ore,  18,000,000  tons  of  fuel  (chiefly  coke),  and  9,500,000  tons  of 
limestone.  They  produced  nearly  18,000,000  tons  of  pig  iron  —  over 
one  third  of  the  world's  output.  Germany  and  the  United  Kingdom 
produced  the  bulk  of  the  remainder. 

Varieties  of  Iron.  —  The  iron  we  use  and  speak  of  is 
not  pure  iron,  but  largely  a  mixture  or  compound  of  iron 
with  other  elements,  chiefly  carbon.  It  is  customary  to 
speak  of  three  varieties  of  iron,  —  cast  iron,  steel,  and 
wrought  iron.  This  classification  is  based  chemically 
upon  the  per  cent  of  carbon  they  contain,  though  their 
physical  properties  are  modified  by  the  presence  of  silicon, 


37 8  Descriptive  Chemistry. 

phosphorus,  sulphur,  and  manganese.  Each  typical  vari- 
ety has  specific  properties ;  but  the  different  varieties  are 
closely  related,  and  pass  easily  and  gradually  into  each 
other.  Commercially,  there  are  several  kinds  of  cast  iron 
and  many  kinds  of  steel. 

Cast  Iron  is  the  most  impure  variety.  It  contains,  be- 
sides carbon,  the  impurities  mentioned  above.  It  has  a  crys- 
talline structure,  and  is  brittle.  The  proportion  of  carbon 
varies  from  1.5  to  6  or  more  per  cent.  If  most  of  the  car- 
bon is  combined  with  the  iron,  the  metal  is  called  white 
cast  iron.  But  if  the  molten  metal  cools  slowly,  much  of 
the  carbon  remains  uncombined  as  graphite,  and  the  color 
of  the  iron  is  gray;  this  kind  is  gray  cast  iron.  It  is 
softer  than  the  white  variety,  and  melts  at  a  lower  tempera- 
ture. Although  cast  iron  is  brittle,  it  will  withstand  great 
pressure.  It  cannot  be  welded  or  forged,  that  is,  hot 
pieces  cannot  be  united,  nor  be  shaped  by  hammering. 
But  it  is  extensively  used  to  make  castings.  This  is  the 
kind  of  iron  used  in  an  ordinary  iron  foundry.  The  iron, 
which  melts  at  a  comparatively  low  temperature  (about 
1100°  C.),  is  heated  in  a  furnace  similar  to  a  blast  furnace, 
and  when  molten  is  poured  into  sand  molds  of  the 
desired  shape.  Stoves,  pipes,  pillars,  railings,  parts  of 
machines,  and  many  other  useful  objects  are  made  of  cast 
iron.  Birmingham,  Alabama,  is  the  center  of  the  cast-iron 
industry  in  the  United  States. 

Cast  iron  containing  much  manganese  is  called  spiegel  iron  or  ferro- 
manganese  (see  Manganese).  About  300,000  tons  of  this  kind  are 
annually  produced  in  the  United  States. 

Wrought  Iron  is  the  purest  variety  of  commercial  iron. 
It  contains  not  more  than  0.5  per  cent  of  carbon  and  some- 
times only  0.06  per  cent,  the  average  being  o.  1 5  per  cent. 
It  is  tough,  malleable,  and  fibrous.  It  can  be.  bent.  Un- 


Iron,  Nickel,  and  Cobalt.  379 

like  cast  iron,  it  does  not  withstand  pressure,  but  it  will 
sustain  great  weight.  An  iron  wire  will  sustain  the  weight 
of  nearly  a  mile  of  itself.  It  melts  at  such  a  high  tem- 
perature (1600°  to  2000°  C.)  that  it  is  not  used  for  casting. 
It  can  be  forged  and  welded,  and  is  therefore  often  called 
malleable  iron.  It  may  be  seen  undergoing  these  opera- 
tions in  a  blacksmith's  shop.  It  can  also  be  rolled  into 
plates  and  sheets  and  drawn  into  fine  wire ;  in  these  forms 
the  metal  is  very  strong.  Wrought  iron  is  made  into  wire, 
sheets,  rods,  nails,  spikes,  bolts,  chains,  anchors,  horseshoes, 
tires,  and  agricultural  implements.  It  is  less  important 
than  formerly,  since  it  is  being  largely  replaced  by  steel. 

Wrought  iron  is  made  from  cast  iron  by  burning  out  the 
impurities.  The  process  is  technically  called  puddling. 
Cast  iron  is  heated  in  a  furnace,  much  like  a  reverberatory 
furnace,  lined  on  the  bottom  and  sides  with  iron  ore  (fer- 
ric oxide,  Fe2O3).  The  intense  heat  melts  the  cast  iron  ; 
its  carbon  and  silicon  are  removed  partly  by  the  oxygen  of 
the  air,  but  mainly  by  the  oxygen  of  the  iron  oxide.  As 
the  mass  becomes  pasty,  owing  to  its  higher  melting  point, 
it  is  stirred  vigorously,  or  "  puddled."  At  the  proper  time 
the  lumps  are  removed  and  hammered,  or  more  often 
rolled  between  ponderous  rollers.  This  operation  removes 
the  slag,  and  if  the  rolling  is  repeated,  the  quality  of  the 
iron  is  improved ;  the  final  rolling  leaves  the  iron  in  the 
shape  desired  for  market. 

Steel  is  intermediate  between  cast  iron  and  wrought  iron 
as  far  as  its  proportion  of  carbon  is  concerned.  Many 
grades  of  steel  are  manufactured,  and  their  physical  prop- 
erties depend  not  only  upon  the  presence  of  other  elements 
besides  carbon,  especially  phosphorus,  silicon,  and  certain 
metals,  but  also  upon  the  raw  materials,  the  method  of 
manufacture,  and  subsequent  treatment. 


380  Descriptive  Chemistry. 

The  properties  of  steel  are  numerous.  It  is  both 
fusible  and  malleable,  and  hence  can  be  forged,  welded, 
and  cast.  It  is  harder,  stronger,  and  more  durable  than 
pure  iron,  and  is  more  serviceable.  But  its  most  valuable 
property  is  the  varying  hardness  which  it  can  be  made  to 
acquire.  If  steel  is  heated  very  hot  and  then  suddenly 
cooled  by  immersion  in  cold  water  or  oil,  it  becomes  brittle 
and  very^hard.  But  if  heated  and  then  cooled  slowly,  it 
becomes  soft,  tough,  and  elastic.  All  grades  of  hardness 
may  be  obtained  between  these  extremes.  And  if  the 
hardened  steel  is  reheated  to  a  definite  temperature,  deter- 
mined by  the  color  the  metal  assumes,  and  then  properly 
cooled,  a  definite  degree  of  hardness  and  elasticity  is  ob- 
tained. This  last  operation  is  called  tempering.  Every 
kind  of  tool  has  a  temper  determined  by  its  use.  Special 
grades  of  hard  steel  are  also  made  by  the  addition  of  cer- 
tain metals,  especially  chromium  and  nickel.  Harveyized 
steel  is  made  by  packing  steel  in  a  mixture  of  charcoal 
and  boneblack,  and  heating  it  to  a  very  high  temperature. 
This  operation  hardens  the  surface.  This  brand  of  steel 
is  extensively  used  as  armor  plate  in  warships. 

Manufacture  of  Steel.  —  The  aim  in  the  manufacture 
of  steel  is  to  prepare  a  product  containing  little  or  no  sul- 
phur, phosphorus,  and  silicon,  but  the  desired  proportion 
of  carbon.  This  may  be  done  by  three  general  methods  : 

(1)  the  carbon,  may  be  partly   removed    from    cast   iron, 

(2)  carbon  may  be  added  to  wrought  iron,  (3)  cast  iron 
may  be  added  to  wrought  iron.     The  first  method  is  diffi- 
cult to  operate,  and  is  seldom  used.     The  other  methods 
are  utilized  by  several  processes. 

(i)  In  the  cementation  or  crucible  process,  wrought 
iron  and  carbon  are  packed  in  tight  fire-clay  boxes  and 
heated  for  several  days,  The  iron  slowly  absorbs  carbon 


Iron,  Nickel,  and  Cobalt. 


381 


in  some  way  unknown  at  present,  and  becomes  a  steel  of 
extreme  purity  and  excellent  quality.  The  bars  are  melted 
in  graphite  crucibles  to  make  the  metal  of  uniform  quality, 
and  cast  into  large  bars  called  ingots.  This  process  is 
long  and  expensive,  but  the  steel  is  considered  the  best  for 
fine  tools. 

(2)  The  Bessemer  process  is  the  one  in  most  general 
use.  It  was  devised  in  about  1860,  and  has  practically 
revolutionized  steel  making.  By  the  economical,  scien- 
tific, and  extensive  application  of  this  process,  all  grades  of 
steel  are  quickly  made  at  such  a  relatively  small  cost  that 
the  use  of  this  metal  has  been  enormously  extended,  much 
to  the  prosperity  of  the  United  States.  About  two  thirds 
of  the  annual  production  is  Bessemer  steel.  The  process 
consists  in  burning  out  the  impurities  in  cast  iron  by  forc- 
ing air  through  the  molten  metal,  and  then  adding  just 
enough  cast  iron  (spiegel  iron)  to  give  the  desired  propor- 
tion of  carbon.  The  operation  is 
carried  on  in  a  converter  (Fig.  73). 
This  is  a  huge,  egg-shaped  vessel, 
supported  so  that  it  can  be  rotated 
into  different  positions;  it  is  also 
provided  with  holes  at  the  bottom 
through  which  a  powerful  blast  of 
air  can  be  blown.  It  is  made  of  thick 
wrought  iron  plates,  and  is  lined  with 
an  infusible  mixture  rich  in  silica. 
The  converter  is  swung  into  a  hori- 


FlG.  73.  —  Converter. 


zontal  position  and  five  to  twenty  tons  of  molten  pig  iron 
are  poured  in  direct  from  the  blast  furnace.  The  air  blast 
is  turned  on,  and  the  converter  is  swung  back  to  a  vertical 
position.  As  the  air  is  forced  through  the  molten  metal, 
the  temperature  rises,  the  carbon  is  oxidized  to  carbon 


382  Descriptive  Chemistry. 

monoxide  which  burns  on  the  surface  of  the  metal,  and  the 
silicon  is  oxidized  to  silicon  dioxide,  which  is  taken  up  by 
the  slag.  This  oxidation  generates  enough  heat  to  keep 
the  metal  melted,  and  no  fuel  need  be  used.  As  soon 
as  the  impurities  have  been  burned  out,  sufficient  spiegel 
iron  is  added  to  change  the  wrought  iron  into  steel.  By 
adding  spiegel  iron  of  known  composition,  Bessemer  steel 
of  any  desired  grade  is  produced.  After  the  completion 
of  the  operation,  which  takes  about  twenty  minutes,  the 
contents  of  the  converter  are  poured  into  molds. 

(3)  In  the  Bessemer  process,  sulphur  and   phosphorus 
are  not  removed.     Both  are  objectionable  impurities ;  sul- 
phur makes  steel  brittle  when  hot,  and  phosphorus  makes 
it  brittle  when  cold.     The  Thomas-Gilchrist  process  is  a 
modification  of  the  Bessemer  process  by  which  the  sulphur 
and  phosphorus  can  be  removed.     The  converter  is  lined 
with  a  mixture  of  lime  and  magnesia,  called  a  basic  lining, 
lime  is  also  added  to  the  charge  of  pig  iron,  and  the  blast 
is  continued  a  little  longer  than  in  the  Bessemer  process, 
otherwise  the  operations  are  the  same.     The  phosphorus 
forms  a  phosphate  and  the   sulphur  a  sulphate,  both  of 
which  are  taken  up  by  the  lining.     The  lining,  which  is 
known  as  Thomas  slag,  is  used  as  a  source  of  phosphorus 
for  fertilizers. 

(4)  In  the  Siemens-Martin  or  open-hearth  process  a 
mixture  of  cast  iron  and  wrought  iron  (or  steel)  in  proper 
proportions  is  melted  on  a  hearth  with  an  oxidizing  gas 
flame.     Old  wrought  iron  or  cast  iron,  known  as  "scrap," 
can  be  used.     When  a  test  shows  that  the  metal  contains 
the  desired  proportion  of  carbon,  ferromanganese  is  added, 
and  the  charge  is  then  poured  into  molds.     This  process 
requires  a  special  furnace  and  gas  plant,  and  is  more  ex- 
pensive than  the  Bessemer  process,  since  it  takes  longer. 


Iron,  Nickel,  and  Cobalt.  383 

But  it  is  easily  controlled,  and  yields  a  tough,  elastic  steel, 
which  is  excellent  for  bridges,  large  machines,  large  guns, 
and  gun  carriages.  Immense,  quantities  of  the  nickel  steel 
used  for  the  armor  plate  are  made  by  the  open-hearth 
process.  The  production  of  the  open-hearth  steel  has  more 
than  doubled  in  the  last  few  years,  being  over  five  and 
a  half  million  tons  in  1902. 

Uses  of  Steel.  —  Steel  is  now  used  instead  of  iron  for 
many  purposes.  High  buildings,  bridges,  rails,  cars,  loco- 
motives, battleships,  electrical  machinery,  boilers,  agricul- 
tural implements,  wire  nails,  rods,  hoops,  tin  plates,  and 
castings  of  all  kinds  consume  vast  amounts.  Its  extensive 
use  in  making  springs,  tools,  cutlery,  pens,  needles,  etc., 
need  not  be  further  mentioned. 

Properties  of  Iron.  —  Chemically  pure  iron,  though  un- 
known in  commerce,  may  be  obtained  in  the  laboratory 
by  reducing  the  oxide  or  chloride  with  hydrogen  or  with 
alcohol.  Such  iron  is  called  iron  "  by  hydrogen,"  or  "  by 
alcohol."  The  purest  commercial  form  is  the  wrought 
iron  used  for  piano  wire.  Pure  iron  is  a  silvery  white, 
lustrous,  metal.  It  is  softer  than  ordinary  iron,  but  melts 
at  a  higher  temperature.  The  specific  gravity  is  about 
7.8.  It  is  attracted  by  a  magnet,  but  soon  loses  its  own 
magnetism.  Dry  air  has  no  effect  upon  iron,  but  moist 
air  containing  carbon  dioxide  rusts  it.  Iron  rust  is  a  com- 
plex compound,  but  its  essential  constituent  is  a  ferric 
hydroxide  (Fe2O3  .  Fe2(OH)6).  Rusting  proceeds  rapidly, 
because  the  film  of  rust  is  not  compact  enough  to  protect 
the  metal.  Like  many  metals,  iron  readily  interacts  with 
dilute  acids,  and  as  a  rule  hydrogen  andv  ferrous  com- 
pounds are  the  products. 


384  Descriptive  Chemistry. 

With  nitric  acid  various  products  result,  according  to  the  conditions, 
—  ferrous  nitrate  and  ammonium  nitrate,  if  the  acid  is  cold,  but  ferric 
nitrate  and  oxides  of  nitrogen  if  the  acid  is  warm.  If  a  clean  iron  wire  is 
dipped  into  fuming  nitric  acid  and  then  into  ordinary  nitric  acid,  no  action 
is  apparent.  The  iron  is  said  to  be  passive.  This  peculiar  fact  has  not 
been  adequately  explained.  Steam  and  hot  iron  interact,  thus  — 

3Fe      +      4H2O      =      Fe3O4      +      4H2 
Iron  Water          Iron  Oxide      Hydrogen 

(See  Preparation  of  Hydrogen.) 

Compounds  of  Iron.  —  Iron  forms  two  series  of  com- 
pounds,—  the  ferrous  and  the  ferric.  They  are  analogous 
to  cuprous  and  cupric,  mercurous  and  mercuric  com- 
pounds. Ferrous  compounds  in  an  acid  solution  pass  into 
the  corresponding  ferric  compound  by  the  action  of  oxi- 
dizing agents,  e.g.  oxygen,  nitric  acid,  potassium  chlorate, 
potassium  permanganate,  and  chlorine.  Conversely,  ferric 
compounds  are  reduced  to  the  ferrous  by  reducing  agents, 
e.g.  hydrogen,  hydrogen  sulphide,  sulphur  dioxide,  and 
stannous  chloride.  The  passage  from  one  series  to  the 
other  occurs  easily,  especially  from  ferrous  to  ferric.  In 
most  of  its  compounds,  iron  acts  as  a  metal.  Many  com- 
pounds of  iron  have  industrial  importance,  as  well  as 
scientific  interest. 

Oxides  and  Hydroxides  of  Iron.  —  Iron  forms  three 
oxides.  Ferrous  oxide  (FeO)  is  an  unstable  black  powder. 
Ferric  oxide  (Fe2O3)  occurs  native  in  many  varieties  as 
hematite  —  the  most  abundant  ore  of  iron.  It  may  be  pre- 
pared by  heating  ferrous  sulphate  or  ferric  hydroxide. 
Large  quantities  are  obtained  as  a  by-product  in  the  manu- 
facture of  Nordhausen  (or  fuming)  sulphuric  acid  and  of 
galvanized  iron  and  tinned  ware.  It  is  sold  under  the 
names  rouge,  crocus,  and  Venetian  red.  It  is  used  to  pol- 
ish glass  and  jewelry,  and  to  make  red  paint.  Ferrous- 


Iron,  Nickel,  and  Cobalt.  385 

ferric  or  ferroso-ferric  oxide  (magnetic  oxide  of  iron, 
Fe3O4)  occurs  native  as  magnetite ;  if  noticeably  magnetic, 
it  is  called  loadstone.  It  is  produced  as  a  black  film  or 
scale  by  heating  iron  in  the  air ;  heaps  of  it  are  often  seen 
beside  the  anvil  in  a  blacksmith's  shop.  The  firm  coating 
of  this  oxide  formed  by  exposing  -iron  to  steam  protects 
the  metal  from  further  oxidation. 

Ferrous  hydroxide  (Fe(OH)2)  is  a  white  solid  formed  by  the  inter- 
action of  a  ferrous  salt  and  an  alkali,  such  as  sodium  hydroxide.  Ex- 
posed to  the  air,  it  soon  turns  green,  and  finally  brown,  owing  to  the 
formation  of  ferric  hydroxide.  Ferric  hydroxide  (Fe2(OH)6)  is  a  red- 
dish brown  solid,  formed  by  the  interaction  of  ammonium  hydroxide 
(or  any  alkali)  and  a  ferric  salt.  Several  ferric  hydroxides  are  known. 
The  freshly  prepared  compound  is  an  antidote  for  arsenic. 

Ferrous  Sulphate  (FeSO4)  is  a  green  salt  obtained  by 
the  interaction  of  iron  (or  ferrous  sulphide)  and  dilute  sul- 
phuric acid,  and  is  a  by-product  in  several  industries  (e.g. 
see  Ferric  Oxide).  It  is  prepared  on  a  large  scale  by  oxi- 
dizing iron  pyrites  (FeS2);  this  is  accomplished  simply  by 
roasting,  or  more  often  by  exposing  heaps  of  pyrites  to 
moist  air.  The  mass  is  extracted  with  water  containing 
scrap  iron  and  a  small  proportion  of  sulphuric  acid.  From 
the  clear  solution,  large  light  green  crystals  are  obtained. 
The  crystallized  salt  (FeSO4 .  7  H2O)  is  also  called  green 
vitriol  or  copperas.  Exposed  to  the  air,  ferrous  sulphate 
effloresces  and  oxidizes.  Large  quantities  are  used  as  a 
mordant  and  a  disinfectant,  and  in  manufacturing  ink, 
bluing,  and  pigments.  Much  black  writing  ink  is  made 
essentially  by  mixing  ferrous  sulphate,  nutgalls,  gum,  and 
water.  Blue  ink  is  usually  made  of  Prussian  blue  —  an 
iron  compound  (see  below)  —  oxalic  acid,  and  water. 

Ferric  Sulphate  (Fe2(SO4)3)  is  formed  by  oxidizing  an  acid  solution 
of  ferrous  sulphate  with  nitric  acid.  When  ferric  sulphate  solution  is 


j 86  Descriptive  Chemistry. 

mixed  with  the  proper  quantity  of  potassium  (or  ammonium)  sulphate, 
iron  alum  (K2Fe2(SO4)4  .  24  H2O)  is  formed.  It  is  a  violet,  crystallized 
solid,  which  has  properties  like  ordinary  alum.  Iron  alum  is  used 
chiefly  as  a  mordant. 

Iron  Sulphides.  —  There  are  two  iron  sulphides.  Com- 
mercial ferrous  sulphide  (FeS)  is  a  black,  brittle,  me- 
tallic-looking solid,  but  the  pure  compound  is  yellow  and 
crystalline.  It  is  also  obtained  as  a  black  powder  by  the 
interaction  of  a  dissolved  ferric  or  ferrous  salt  and  ammo- 
nium (or  potassium)  sulphide.  It  is  made  on  a  large  scale 
by  fusing  a  mixture  of  iron  and  sulphur.  It  is  used  chiefly 
in  preparing  hydrogen  sulphide.  Ferric  sulphide  (iron 
disulphide,  iron  pyrites,  pyrite,  FeS2)  is  one  of  the  com- 
monest minerals.  It  is  a  lustrous,  metallic,  brass-yellow 
solid.  Crystals  of  pyrites,  found  in  many  rocks,  are  often 
mistaken  for  gold — hence  the  popular  name  "fool's  gold." 
It  is  valueless  as  an  iron  ore,  but  large  quantities  are  used 
as  a  source  of  sulphur  in  making  sulphuric  acid.  Over  one 
and  a  half  million  tons  are  annually  consumed  in  the  acid 
industry.  The  largest  pyrite  producers  are  Spain,  France, 
Portugal,  Germany,  and  the  United  States.  The  domestic 
output  comes  chiefly  from  Virginia,  Colorado,  Massachu- 
setts, and  New  York. 

Iron  Chlorides.  —  When  iron  interacts  with  hydrochloric  acid,  fer- 
rous chloride  (FeCl2)  is  formed  in  solution.  Heated  in  the  air,  or 
better  with  potassium  chlorate  or  nitric  acid,  it  is  changed  into  ferric 
chloride,  thus  — 

2FeCl2       +         2HC1        +       O       =     2FeCl3       +    H2O 
Ferrous  Chlo-       Hydrochloric      Oxygen      Ferric  Chlo-      Water 
ride  Acid  ride 

Ferric  chloride  is  a  black,  lustrous,  crystalline  solid ;  but  owing  to  its 
extreme  deliquescence,  it  is  usually  sold  as  a  solution,  which  is  a  dark 
brown  liquid.  It  is  prepared  by  passing  chlorine  into  a  ferrous  chloride 


Iron,  Nickel,  and  Cobalt.  387 

solution,  or  by  the  interaction  of  iron  and  aqua  regia.  When  treated 
with  nascent  hydrogen  or  another  reducing  agent,  ferric  chloride  is 
changed  into  ferrous  chloride. 

Ferrous  Carbonate  (FeCO3)  occurs  native  as  the  iron  ore  siderite, 
clay  iron  stone,  or  spathic  iron  ore.  The  typical  variety  is  light  yellow 
or  brown,  lustrous,  crystalline,  and  not  very  hard ;  but  many  kinds  are 
impure,  and  the  properties  vary.  It  is  slightly  soluble  in  water  contain- 
ing carbon  dioxide,  and  is  therefore  found  in  some  mineral  springs  (see 
Chalybeate  Waters).  Like  all  carbonates,  it  yields  carbon  dioxide 
with  warm  hydrochloric  acid. 

Iron  Cyanides.  —  Iron  and  cyanogen  (CN),  with  or  with- 
out potassium,  form  several  compounds.  The  most  impor- 
tant is  potassium  ferrocyanide  (K4Fe(CN)6).  It  is  a 
lemon-yellow,  crystallized  solid,  containing  three  molecules 
of  water  of  crystallization.  Unlike  most  cyanogen  com- 
pounds, it  is  not  poisonous.  Its  commercial  name  is 
yellow  prussiate  of  potash.  It  is  manufactured  by  fusing 
together  iron  filings,  potassium  carbonate,  and  nitrogenous 
animal  matter  (such  as  horn,  hair,  blood,  feathers,  and 
leather).  The  mass  is  extracted  with  water,  and  the  salt 
is  separated  by  crystallization.  In  Germany  this  salt  is 
manufactured  from  the  iron  oxide  which  has  been  used 
to  purify  illuminating  gas.  Large  quantities  are  used  in 
dyeing  and  calico  printing,  and  in  making  bluing  and 
potassium  cyanogen  compounds.  Potassium  ferricyanide 
(K3Fe(CN)6)  is  a  dark  red,  crystalline  solid,  containing  no 
water  of  crystallization.  It  is  often  called  red  prussiate 
of  potash.  It  is  manufactured  by  oxidizing  potassium 
ferrocyanide  with  chlorine,  thus  — 

K4Fe(CN)6     +       Cl      =     K3Fe(CN)6    +      KC1 
Potassium  Ferro-        Chlorine        Potassium  Ferri-        Potassium 
cyanide  cyanide  Chloride 

It  is  very  soluble  in  water,  forming  a  deep  yellow,  unstable 
solution.  In  alkaline  solution  it  is  a  vigorous  oxidizing 


388  Descriptive  Chemistry. 

agent,  and  therefore  finds  extensive  use  in  dyeing.  It  is 
also  used  as  one  of  the  ingredients  of  the  sensitive  coating 
of  "  blue  print  "  paper. 

Ferrous  salts  and  potassium  ferricyanide  interact  in  solution  and  pro- 
duce ferrous  ferricyanide  (Fe3(Fe(CN)6)2) .  This  is  a  blue  solid  and  is 
often  called  TurnbulPs  blue.  But  ferrous  salts  produce  with  potassium 
ferrocyanide  a  white  precipitate  (ferrous  ferrocyanide),  which  quickly  oxi- 
dizes to  a  complex  blue  compound.  Ferric  salts  interact  with  potassium 
ferrocyanide  and  produce  ferric  ferrocyanide  (Fe4(Fe(CN)6)2).  This 
is  a  dark  blue  solid,  and  is  called  Prussian  blue  or  Berlin  blue.  Ferric 
salts  produce  no  precipitate  with  potassium  ferricyanide.  Prussian  blue 
is  extensively  used  in  dyeing  and  calico  printing,  and  in  making  bluing. 
The  above  reactions,  which  allow  ferrous  and  ferric  salts  to  be  distin- 
guished, may  be  summarized  as  follows  :  — 


CYANIDE. 

FERROUS  SALT. 

FERRIC  SALT. 

Ferrocyanide 
Ferricyanide 

Whitish  precipitate 
Turnbuirs  blue 

Prussian  blue 
No  precipitate 

Besides  the  above  tests,  potassium  sulphocyanate  produces  a  dark  red 
liquid  with  ferric  salts,  but  leaves  ferrous  salts  unchanged.  The  tests 
for  iron  are  thus  numerous  and  specific. 

NICKEL. 

Nickel,  Ni,  occurs  combined  with  arsenic,  sulphur,  or 
both.  Small  amounts  of  metallic  nickel  are  found  in  me- 
teorites. The  chief  ores  are  nickel-bearing  iron  sulphides, 
which  are  abundant  in  the  Sudbury  district,  Canada,  and 
the  silicates  found  in  New  Caledonia.  A  small  amount  is 
produced  in  the  United  States  as  a  by-product  in  smelt- 
ing lead  ores  from  a  Missouri  mine. 

Nickel  is  obtained  from  its  Ores  by  complicated  pro- 
cesses, and  is  now  refined  by  electrolysis.  It  is  a  white 


Iron,  Nickel,  and  Cobalt.  389 

metal,  which  takes  a  brilliant  polish.  It  is  ductile,  hard, 
tenacious,  and  does  not  tarnish  in  the  air.  Like  cobalt,  it 
is  attracted  by  a  magnet. 

Nickel  has  varied  Uses.  —  For  many  years  it  has  been 
used  as  one  ingredient  of  the  small  coins  of  several  coun- 
tries. The  per  cent  of  nickel  varies  from  12  in  the  United 
States  cent  to  25  in  the  five-cent  piece.  German  silver 
contains  from  15  to  25  per  cent  of  nickel,  the  rest  being 
copper  and  zinc.  Large  quantities  of  nickel  are  used  to 
coat  or  plate  other  metals,  especially  iron  and  brass.  The 
nickel  plating  is  done  by  electrolysis,  as  in  the  case  of 
silver  and  gold  plating,  though  the  electrolytic  solution 
used  is  a  sulphate  of  nickel  and  ammonium,  not  a  cyanide. 
The  deposit  of  nickel  is  hard,  brilliant,  and  durable. 
Nickel  becomes  malleable,  if  a  little  magnesium  is  added 
to  the  molten  metal,  and  sheets  of  iron  covered  with  such 
nickel  are  made  into  vessels  for  cooking.  Nickeloid  is  a 
nickel-plated  sheet  zinc.  Its  attractive  appearance  and 
non-corrosive  property  adapt  it  for  the  manufacture  of 
reflectors,  refrigerator  linings,  bath  tubs,  show  cases,  and 
signs.  The  most  important  use  of  nickel  is  in  the  manu- 
facture of  nickel  steel.  This  contains  about  3.5  per  cent 
of  nickel.  Large  quantities  are  used  for  the  armor  plates 
and  turrets  of  battleships,,  and  for  parts  of  machinery 
requiring  great  strength. 

Nickel  forms  two  series  of  compounds,  —  the  nickelous  and  the  nick- 
elic.  The  nickelous  are  more  common,  and  many  of  them  are  green. 
The  test  for  nickel  is  the  formation  of  the  apple-green  hydroxide 
(Ni(OH)2)  by  the  interaction  of  an  alkali  and  the  solution  of  a  nickel 
salt. 

Cobalt,  Co,  generally  occurs  combined  with  arsenic  and  sulphur,  and 
is  often  associated  with  nickel  compounds.  It  is  a  lustrous  metal  with 
a  reddish  tinge,  harder  than  iron,  but  less  magnetic.  The  hydrated 


390  Descriptive  Chemistry. 

compounds  are  red  in  solution,  anhydrous  compounds  are  blue. 
Hence  red  crystallized  salts  turn  blue  when  heated.  Some  cobalt 
compounds  are  used  to  color  glass,  porcelain,  and  paper,  especially  a 
cobalt  silicate.  This  is  known  as  smalt,  or  smalt  blue ;  and  since  it  is 
unchanged  by  sunlight,  acids,  or  alkalies,  it  is  used  to  decorate  porce- 
lain. Other  pigments  are  cobalt  blue  (an  oxide  of  cobalt  and  aluminium) , 
and  Rinmann's  green  (an  oxide  of  cobalt  and  zinc).  The  blue  color 
produced  by  fusing  cobalt  compounds  into  a  borax  bead  is  the  test  for 
cobalt. 

EXERCISES. 

1 .  What  is  the  symbol  of  iron  ?     From  what  word  is  it  derived  ? 

2.  Discuss  the  occurrence  of  iron.     Name  the  chief  ores.     Name 
other  compounds  of  iron.     What  proportion  of  the  earth's  crust  is  iron? 

3.  Discuss  («)  the  production  and  transportation  of  iron  ore   in 
the  United  States,  and  ($)  the  production  of  iron. 

4.  What   is   the   general  chemical   change    in    the    metallurgy  of 
iron?     Describe  a  blast    furnace.     Summarize    the    smelting   of   iron. 
Discuss  the  chief  physical  and  chemical  changes  involved  in  the  smelting. 

5.  Name  the  varieties  of  iron.     How   do  they  differ  essentially? 
What  is  (a)  galvanized  iron,  (b)  meteoric  iron? 

6.  Describe  cast  iron.     State  its  composition,  properties,  and  uses. 

7.  Describe  the  manufacture  of  wrought  iron.     State  its  composi- 
tion, properties,  and  uses. 

8.  State  the  composition  and  properties  of  steel.     Compare  briefly 
with  cast  and  wrought  iron.     What  is  tempering? 

9.  Describe  the  manufacture  of  steel  by  the  following  processes : 
(a)  cementation,  (b}  Bessemer,  (c)  Thomas-Gilchrist,  (d}  Siemens- 
Martin. 

10.  State  the  uses  of  steel. 

1 1 .  State  the  properties  of  iron. 

12.  How  are  ferrous  changed  into  ferric  compounds,  and  •vice  versa  ? 

13.  How  is  ferric  oxide  prepared?     What  is  the  native  form  called? 
For  what  is  crocus  used? 

14.  What  is  the  formula  and  chemical  name  of  magnetic  oxide  of 
iron?     What  is  loadstone?     How  is  magnetic  oxide  of  iron  produced? 
What  is  the  native  form  called? 

15.  Describe  ferric  hydroxide.     What  is  its  use? 

16.  Describe  ferrous  sulphate.     How  is  it  prepared?    For  what  is 
it  used  ?    What  is  copperas  ? 


Iron,  Nickel,  and  Cobalt.  391 

17.  What  is  iron  alum  ?     How  is  it  related  to  ordinary  alum? 

1 8.  Describe  ferrous  sulphide.     How  is  it  made?     For  what  is  it 
used?     Compare  it  with  ferric  sulphide.     Discuss  the  occurrence  and 
use  of  the  latter. 

19.  Describe  ferrous  carbonate. 

20.  Describe  potassium  ferrocyanide.     How  is  it  made?     State  its 
properties  and  uses.     What  is  its  common  name?     Its  formula? 

21.  Describe  potassium  ferricyanide.     For  what  is  it  used?     How 
is  it  related  chemically  to  potassium  ferrocyanide? 

22.  Describe  the  tests  for  iron.     What  is  Prussian  blue  ?     For  what 
is  it  used? 

23.  Discuss  the  occurrence  of  nickel.     State  its  properties  and  uses. 
Describe  nickel  plating.     What  is   (a)  a  "nickel,"  (b)   nickel  steel? 
What  is  the  test  for  nickel  ? 

24.  State  the    properties  of  cobalt.     For  what  are  its  compounds 
used?    What  is  smalt?     What  is  the  test  for  cobalt ? 

PROBLEMS. 

1.  Calculate  the  percentage  composition  of  («)  ferric  oxide,  (b)  fer- 
rous sulphate,  (V)  ferrous  sulphide  (FeS). 

2.  If  1.586  gm.  of  iron  form  2.265  Km-  °f  ferric  oxide,  what  is 
the  atomic  weight  of  iron?     (Equation  is  2  Fe  +  3  O  =  Fe2O3.) 


CHAPTER  XXIX. 
PLATINUM  AND  ASSOCIATED  METALS. 

Occurrence  of  Platinum.  —  Platinum  occurs  as  the  essen- 
tial ingredient  of  platinum  ore  or  so-called  native  platinum. 
The  ore  contains  from  60  to  86  per  cent  of  platinum.  The 
other  metals  present  are  ruthenium,  osmium,  iridium,  rho- 
dium, and  palladium.  Iron,  gold,  and  copper  are  also  usu- 
ally present.  Only  one  native  compound  is  known,  viz. 
platinum  arsenide  (sperrylite,  PtAs2). 

The  ore  is  found  chiefly  in  the  Ural  Mountains  in  Russia,  but  some 
comes  from  South  America,  Australia,  and  Borneo.  The  United  States 
produced  about  1400  ounces  of  metallic  platinum  in  1901  — the  largest 
annual  output  on  record.  It  came  from  the  gold  deposits  in  California 
and  the  copper  mines  in  Wyoming.  The  latter  source  also  furnished 
osmium,  palladium,  and  iridium.  The  world's  annual  production  of 
metallic  platinum  for  the  last  few  years  has  been  about  165,000  ounces. 
Russia  supplies  over  90  per  cent  of  this  amount. 

The  word  platinum  is  derived  from  platina,  a  form  of  the  Spanish 
word  plat  a,  meaning  silver,  because  native  platinum  was  regarded  as  an 
impure  ore  of  silver  by  the  Spaniards,  who  first  discovered  it  in  South 
America  about  1735.  Platinum  is  now  sometimes  called  by  its  old 
name  platina. 

Preparation  of  Platinum.  —  The  platinum  ore,  which  occurs  in 
rounded  grains  or  flattened  scales,  is  first  digested  with  dilute  aqua 
regia  to  remove  the  gold,  silver,  and  copper ;  and  then  with  concen- 
trated aqua  regia,  which  changes  all  the  platinum  and  a  very  little 
iridium  into  soluble  compounds,  leaving  behind  an  alloy  of  iridium  and 
osmium.  From  the  clear  solution  the  platinum  and  iridium  are  precipi- 
tated by  ammonium  chloride  as  compounds,  which,  on  heating,  yield 
the  metals  as  a  spongy  mass.  This  spongy  platinum  is  melted  in  a 

392 


Platinum  and  Associated  Metals.  393 

lime  crucible  with  an  oxhydrogen  flame,  or  hammered  while  hot  into 
sheet  platinum.  The  very  small  amount  of  iridium  is  seldom  removed 
from  the  metallic  platinum. 

Properties  and  Uses  of  Platinum.  —  Platinum  is  a  lus- 
trous, grayish  white  metal.  It  is  malleable  and  ductile, 
and  usually  appears  in  commerce  in  the  form  of  wire  and 
sheets.  Sheet  platinum  is  cut  into  squares — the  familiar 
platinum  foil  of  the  laboratory,  or  made  into  crucibles, 
dishes,  Vid  stills  for  sulphuric  and  hydrofluoric  acid  (Fig. 
74).  Its  use  in  these  forms  is  due  partly  to  its  infusibility 
and  partly  to  its  resistance  to  acids  and  other  corrosive 
chemicals.  Although  it  is  attacked  by  fused  caustic  alka- 
lies and  a  few 
other  substances, 
it  is  practically 
indispensable  in 
the  chemical  lab- 
oratory. Plati- 
num is  a  good 

FIG.  74. —  A  platinum  dish. 

conductor  or  elec- 
tricity, and  large  quantities  are  consumed  in  incandescent 
electric  light  bulbs.  Short  pieces  of  wire  are  fused 
into  the  glass  at  the  base  of  the  bulb  and  attached  to. 
the  outside  wires,  conveying  the  current  to  and  from  the 
carbon  filament  within.  Platinum  is  the  only  metal  thus 
far  found  which  is  perfectly  adapted  to  this  use.  Dentists 
use  alloys  of  platinum  as  a  filling  for  teeth,  and  some  is 
made  into  jewelry.  The  demand  exceeds  the  supply,  and 
in  the  last  five  years  the  price  of  this  rare  metal  has 
doubled,  being  $21  an  ounce  in  1902.  Platinum  has  a 
specific  gravity  of  about  21,  which  is  higher  than  that  of 
any  known  substance,  except  osmium  and  iridium.  In  the 
form  of  a  black,  porous  mass  it  is  called  spongy  platinum, 


394  Descriptive  Chemistry. 

and  a  still  finer  form  is  called  platinum  black.  Both  forms 
absorb  large  volumes  of  gases  ;  and  if  a  current  of  the  gas 
is  directed  against  the  metal,  the  gas  often  takes  fire.  Me- 
tallic platinum  has  the  same  property  to  a  less  degree,  for 
it  becomes  red-hot  if  held  in  a  stream  of  illuminating  gas, 
and  often  ignites  the  gas.  Palladium  has  similar  proper- 
ties (see  Occlusion).  Platinum  forms  alloys  with  other 
metals,  and  should  never  be  heated  with  lead,  similar  met- 
als, or  their  compounds,  since  the  alloys  have  a  low  melt- 
ing point.  With  iridium,  however,  it  forms  a  very  hard 
alloy  of  which  the  international  metric  apparatus  is  made. 

Platinic  Chloride  (PtCl4)  is  the  only  important  compound  of  plati- 
num. It  is  a  brownish  solid  formed  by  treating  platinum  with  aqua 
regia  and  evaporating  the  solution  to  dry  ness.  The  solution  is  used  in 
chemical  analysis,  and  in  photography  to  produce  "  platinum  prints.1' 
Chloroplatinic  acid  (H2PtCl6)  forms  complex  salts,  of  which  the  yel- 
low, crystalline  potassium  chlorplatinate  (K2PtCl6)  and  ammonium 
chlorplatinate  ((NH4)2PtCl6)  are  the  best  known. 

The  Metals  associated  with  Platinum  have  limited  uses.  Pal- 
ladium is  used  in  chemical  analysis  to  absorb  hydrogen,  osmium  is 
utilized  in  the  Auer  incandescent  electric  light,  and  a  native  (as  well  as 
an  artificial)  alloy  of  iridium  and  osmium,  called  iridosmine,  is  used  to 
tip  gold  pens. 

EXERCISES. 

i.    Name  the  metals  related  to  platinum. 
2..   Discuss  the  occurrence  of  platinum. 

3.  What  is  (a)  native  platinum,  (b)  spongy  platinum,  (c)  platinum 
black,  (d)  platinum  foil,  (e)  sheet  platinum  ? 

4.  Discuss  the  production  of  platinum. 

5.  What  is  the  symbol  of  platinum  ?     What  is  the  derivation  of  the 
word  platinum  ? 

6.  Describe  the  preparation  of  platinum.     Summarize  its  properties. 
State  its  uses. 

7.  Describe  platinic  chloride. 

8.  State  the  uses  of  the  metals  related  to  platinum. 


Platinum  and  Associated  Metals.  395 

PROBLEMS. 

1.  A  piece  of  platinum  foil  measuring  10.5  cm.  by  1.5  cm.  weighs 
0.723  gm.     Into  how  many  pieces,  each  weighing  i   dg.,  may  it  be 
divided  ? 

2.  The  specific  heat  of  platinum  is  0.0324.     According  to  analysis, 
35.5  gm.  of  chlorine  unite  with  48.6  gm.  of  platinum  to  form  platinic 
chloride.     What  is  (a)  the  atomic  weight  of  platinum,  and   (b)   the 
formula  of  platinic  chloride  ? 


CHAPTER   XXX. 
GENERAL  RELATIONS  OF  THE   ELEMENTS. 

Introduction.  —  In  the  preceding  chapters  emphasis  has 
been  laid  on  individual  elements.  Certain  group  relations 
were  also  pointed  out,  but  little  or  nothing  was  said  con- 
cerning the  elements  as  a  single  large  group.  The  ele- 
ments are  not  independent.  They  possess  certain  funda- 
mental properties,  which  show  that  although  apparently 
very  different,  they  are  really  closely  related.  In  this 
chapter  we  shall  consider  two  topics  which  illustrate  this 
general  fundamental  relationship,  viz.  the  periodic  law  and 
spectrum  analysis. 

THE    PERIODIC    LAW. 

Classification  of  the  Elements.  —  As  the  number  of 
elements  increased,  attempts  were  made  to  classify  them. 
About  the  time  of  Lavoisier  (1743-1794)  they  were  roughly 
divided  into  metals  and  non-metals.  Those  elements 
were  called  metals  which  were  hard,  lustrous,  heavy,  and 
good  conductors  of  heat,  while  the  others  were  called  non- 
metals.  This  classification  proved  to  be  misleading  as 
additional  elements  were  discovered.  It  is  used,  however, 
even  now,  because  many  common  elements  fall  readily  into 
one  of  these  classes. 

Classification  according  to  acid  and  basic  properties 
prevailed  for  a  time.  But  it  was  abandoned  largely  be- 
cause such  a  basis  of  division  excluded  elements  exhibiting 

396 


General   Relations  of  the  Elements.          397 

both  acid  and  basic  properties,  such  as  arsenic,  antimony, 
chromium,  and  aluminium. 

The  elements  have  also  been  classified  according  to  their 
valence  into  six  or  seven  groups  (the  mono-,  di-,  tri-,  etc.). 
But  this  plan  has  been  largely  given  up  on  account  of  so 
many  troublesome  cases  of  variable  and  unsatisfied  valence 
(see  Valence). 

About  1828  Dumas  pointed  out  striking  resemblances 
between  certain  elements,  and  he  suggested  several  groups 
or  families.  For  example:  — 


(I) 

(2) 

(3) 

(4) 

Lithium 
Sodium 
Potassium 

Selenium 
Sulphur 
Oxygen 

Calcium 
Strontium 
Barium 

Nitrogen 
Phosphorus 
Arsenic 

This  classification  was  arbitrarily  based  on  selected  physi- 
cal and  chemical  properties.  It  was  interesting  but  incom- 
plete, because  it  emphasized  resemblances  and  overlooked 
differences  —  that  is,  the  basis  of  comparison  was  not 
broad  enough. 

The  first  actual  progress  began  to  be  made  about  1850, 
when  chemists  became  deeply  interested  in  the  significance 
of  atomic  weights.  Dumas  (in  1857)  and  others  pointed 
out  certain  remarkable  numerical  relations  existing  be- 
tween the  atomic  weights  of  related  elements.  Thus,  the 
atomic  weight  of  sodium  is  half  the  sijm  of  the  atomic 
weights  of  lithium  and  potassium  — 

Li  =  7,  Na  =  23>K  =  39.         ^~  =  23- 
The  same  is  true  of  phosphorus,  arsenic,  and  antimony  — 
P=3i,  As=75>Sb=i20.     3I  +  I2°=7S.5. 


398  Descriptive  Chemistry. 

The  existence  of  other  relations  similar  to  these,  together 
with  a  deep  desire  to  obtain  more  accurate  atomic  weights 
and  a  growing  interest  in  the  properties  of  the  elements 
themselves,  focused  the  attention  of  chemists  at  this  time 
(1855-1865)  upon  the  relation  of  properties  to  atomic 
weights.  Several  things  fostered  the  above  principle. 
One  was  the  atomic  weight  determinations  of  Stas,  whose 
masterly  work  proved  beyond  doubt  that  Prout  was  incor- 
rect when  he  insisted  in  1815  that  the  atomic  weights  are 
whole  numbers.  Another  was  the  acceptance  by  most 
chemists  of  the  same  table  of  atomic  weights.  A  third 
was  the  rapid  accumulation  of  many  facts  about  the  ele- 
ments and  their  compounds.  Chemists  were  ready  for  a 
new  classification  of  the  elements. 

The  Periodic  Classification.  —  Previous  to  1869  no 
classification  included  all  the  elements.  In  that  year  the 
Russian  chemist  Mendeleeff  published  a  classification  of 
the  elements  according  to  the  periodic  law.  His  views 
had  been  partially  anticipated  by  several  chemists,  and 
were  soon  amplified  by  the  German  chemist,  Lothar  Meyer. 
Their  classification  of  the  elements  revealed  a  new  relation 
between  the  properties  of  the  elements  and  their  atomic 
weights.  If  all  the  elements  are  arranged  in  the  order  of 
their  increasing  atomic  weights  beginning  with  lithium, 
their  properties  will  vary  periodically,  i.e.  at  certain  regu- 
lar intervals  or  periods  elements  will  be  found  which  have 
similar  properties.  In  other  words,  a  certain  increase  in 
atomic  weight  causes  a  reappearance  or  return  of  prop- 
erties. The  general  relation  is  often  summarized  in  the 
Periodic  Law  - 

The  properties  of  the  elements  are  periodic  functions  of 
their  atomic  weights. 


General   Relations  of  the   Elements. 


399 


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400  Descriptive  Chemistry. 

Function  here  means  the  exhibition  of  some  special  rela- 
tion, viz.  that  of  properties  to  atomic  weight.  Interpreted 
freely,  the  law  means  (i)  properties  and  atomic  weight  are 
related,  they  depend  upon  each  other;  and  (2)  this  relation 
is  exhibited  again  and  again  as  we  reach  elements  with 
increasing  atomic  weights  at  regular  intervals  in  the  suc- 
cessive arrangement 

The  Periodic  Table  originally  proposed  by  Mendeleeff 
has  been  modified  from  time  to  time,  as  new  facts  have 
necessitated.  The  table  generally  accepted  at  the  present 
time  is  given  on  page  399. 

From  the  table  it  is  seen  that  the  elements  fall  naturally 
into  two  subdivisions,  (i)  Those  in  the  same  vertical  col- 
umn belong  to  the  same  natural  group  or  family.  Thus, 
in  Group  I  are  found  the  alkali  metals,  in  Group  II  the 
alkaline  earth  metals,  in  Group  VII  the  halogens.  (2) 
The  elements  in  the  same  horizontal  row  belong  to  the 
same  period.  The  periodic  variation  of  their  properties  is 
well  illustrated  by  the  second  and  third  periods.  Begin- 
ning with  lithium,  the  general  chemical  properties  vary 
regularly  with  increasing  atomic  weight  Thus,  the  metal- 
lic character  gradually  diminishes  until  fluorine  is  passed 
and  sodium  is  reached;  here  it  reappears.  Proceeding 
onward  from  sodium,  the  same  gradation  of  properties  is 
noticed  until  potassium  is  reached,  and  here  again  the 
marked  metallic  character  in  the  same  way  reappears. 
There  is  no  sudden  change  in  properties  until  we  pass 
from  one  period  to  the  next.  Thus,  fluorine  at  the  end  of 
the  second  period  forms  a  powerful  acid,  but  sodium  at  the 
beginning  of  the  third  period  forms  a  strong  base.  Simi- 
larly, chlorine  is  strongly  acidic ;  but  potassium,  which 
begins  the  next  period,  is  markedly  basic;  chlorine  is  a 
typical  non-metal,  while  potassium  is  a  typical  metal.  Not 


General   Relations  of  the  Elements.          401 

all  elements  fit  the  periodic  classification  equally  well,  but 
the  arrangement  is  at  least  very  suggestive,  and  doubtless 
expresses  an  approximately  truthful  relation. 

The  Gaps  in  the  Periodic  Classification  probably  corre- 
spond to  elements  not  yet  discovered.  Three  such  gaps, 
which  were  in  the  original  table,  have  been  filled.  When 
Mendeleeff  proposed  his  arrangement,  he  predicted  the 
discovery  of  three  elements  having  definite  properties. 
These  elements,  —  gallium,  scandium,  and  germanium, — 
have  since  been  discovered  and  now  occupy  their  pre- 
dicted place  in  the  table.  Possibly  other  gaps  will  be  filled 
by  newly  discovered  elements. 

The  discovery  of  the  predicted  elements  was  not  the  only  immediate 
service  of  MendeleefFs  table.  It  also  emphasized  the  necessity  of  more 
accurate  atomic  weights.  Several  elements  did  not  fall  into  their  proper 
places,  and  careful  investigation  showed  that  their  accepted  atomic 
weights  were  incorrect.  Thus,  the  atomic  weights  of  beryllium  and  in- 
dium were  changed  to  their  present  values,  and  the  present  order  of  the 
platinum  metals  was  adopted ;  cobalt  and  nickel  are  still  being  studied. 
The  position  of  argon,  helium,  and  very  rare  metals  is  still  doubtful, 
owing  to  a  limited  knowledge  of  their  properties  and  atomic  weights. 
Hydrogen,  also,  still  lacks  a  place. 

SPECTRUM    ANALYSIS. 

Introduction.  —  When  light  from  an  ordinary  gas  flame, 
glowing  lime  or  other  solid,  or  a  Welsbach  flame  is  passed 
through  a  prism  and  falls  upon  a  white  surface,  a  long 
band  of  color  is  produced.  The  colors  are  perfectly 
blended,  and  are  arranged  like  the  familiar  colors  of  the 
rainbow.  This  band  of  colors  is  called  a  spectrum.  The 
white  light  has  been  separated  or  analyzed  into  the  col- 
ored. The  examination  and  study  of  the  spectrum  of  a 
substance  is  spectrum  analysis,  and  it  is  accomplished  by 
a  spectroscope. 


402  Descriptive  Chemistry. 

The  Spectroscope  consists  essentially  of  a  prism  and  tubes,  one  of 
which  is  a  telescope  (Fig.  75) .     The  light  enters  a  slit  in  the  tube,  passes 


FIG.  75.  —  A  spectroscope. 

through,  and  falls  upon  the  prism.  Here  it  is  bent  from  its  path,  and  as 
it  emerges  from  the  prism,  it  may  be  viewed  through  the  telescope  as  a 
magnified  spectrum. 

Kinds  of  Spectra.  —  (i)  The  spectrum  of  an  incan- 
descent solid  is  a  continuous  band  of  colors.  (2)  But  the 
spectra  of  gases  are  narrow,  colored,  vertical  bars  or  lines, 
separated  by  black  spaces.  Thus,  sodium  vapor  has  a  yel- 
low line,  potassium  a  red  and  a  violet  line,  and  barium  sev- 
eral lines  where  the  green  and  yellow  parts  of  the  ordinary 
spectrum  occur.  Each  element  which  is  a  gas,  or  which  can 
be  vaporized,  has  its  own  bright  line  spectrum.  The  lines 
always  occupy  the  same  relative  positions,  which  in  most 
cases  have  been  very  carefully  determined.  Therefore, 
when  examined  through  a  spectroscope,  the  yellow  line  of 
sodium  will  always  be  seen  in  its  proper  place,  and  the  red 
and  violet  potassium  lines  in  their  places.  Therefore,  by 
examining  the  light  from  different  substances,  it  is  possi- 


General  Relations  of  the  Elements.         403 

ble  to  tell  what  elements  they  contain.  (3)  The  spectrum 
of  sunlight  is  the  familiar  band  of  colors,  but  it  is  crossed 
vertically  by  many  black  lines,  which  have  fixed  positions 
(Fig.  76).  It  is  believed  that  the  sun  is  a  glowing  hot 
solid,  surrounded  by  very  hot  gases.  It  therefore  should 

AaBC     D      Eb  F         G  If 


111  ill  I 


Hed      Orange  Fellow  Green    Blue          Indigo  Violet 

FIG.  76.  —  Spectrum  of  sunlight  showing  some  of  the  vertical  lines. 

give  the  two  kinds  of  spectra,  —  the  continuous  and  the 
bright  line.  Now  it  has  been  proved  that  the  vapor  of 
an  element  absorbs  the  light  given  out  by  the  same  ele- 
ment when  solid.  Hence  the  dark  lines  which  appear  in 
the  solar  spectrum  are  caused  by  the  absorptive  power  of 
the  gases  in  the  sun's  atmosphere.  The  solar  spectrum  is 
often  called  an  absorptive  spectrum. 

Spectrum  Analysis. — In  the  laboratory  the  spectro- 
scope is  used  to  detect  the  presence  of  certain  elements, 
more  especially  the  metals.  If  the  metal  or  one  of  its 
compounds  is  put  on  a  platinum  wire  and  held  in  the 
Bunsen  flame  before  the  slit,  the  characteristic  spectrum 
of  the  element  can  be  easily  recognized  in  the  telescope. 
Two  spectra  do  not  interfere,  because  each  line  has  its 
own  place.  Hence  several  elements  may  be  distinguished 
in  a  mixture.  Minute  quantities  are  easily  detected  by  the 
spectroscope.  Rare  elements,  which  can  be  obtained  only 
in  very  small  quantities  or  with  great  difficulty,  are  studied 
by  the  spectroscope.  Thus,  Bunsen,  who  (with  Kirch- 
hoff)  devised  the  improved  spectroscope,  discovered  the 
rare  metals,  rubidium  and  caesium.  And  within  the  last 
few  years  the  spectroscope  has  been  especially  serviceable 


404  Descriptive  Chemistry. 

in  studying  argon,  helium,  krypton,  neon,  and  xenon. 
By  means  of  the  spectroscope  it  has  been  shown  that  the 
sun  contains  many  elements  found  in  our  earth.  Accord- 
ing to  a  reliable  authority,  about  thirty  of  the  elements 
known  to  us  are  present  in  the  sun.  The  spectroscope 
also  enables  astronomers  to  tell  the  nature  of  stars,  comets, 
nebulae,  and  other  heavenly  bodies.  The  stars  thus  far 
examined  give  spectra  crossed  by  dark  lines,  and  therefore 
these  bodies  are  like  the  sun ;  but  nebulae  give  bright  line 
spectra,  and  hence  consist  of  incandescent  gases. 

EXERCISES. 

1.  Discuss  the  classification  of  the  elements  according  to  (a)  metals 
and  non-metals,  (£)  acid  and  basic  properties,  (c)  valence,  (d)  groups 
based  on  resemblances,  (e)  numerical  relations. 

2.  What  is   the  fundamental  idea  of  the  periodic  classification  ? 
How  does  it  differ  from  previous  systems  ?     When  and  by  whom  was 
this  classification  proposed  and  developed  ? 

3.  State  the  periodic  law.     Explain  it.    What  is  meant  by  (a)  func- 
tion, (£)  period,  (c}  group  ? 

4.  Illustrate  the  law  by  (a)  the  alkali  metals,  and  (b}  the  halogens. 

5.  Discuss  the  gaps  in  the  periodic  arrangement  of  the  elements. 

6.  Of  what  use  has  this  law  been  ? 

7.  State  some  objections  to  it. 

8.  Describe  (a)  a  continuous  spectrum,  (£)  a  line  spectrum,  (c)  an 
absorption  spectrum. 

9.  Describe  a  spectroscope.     How  is  it  used  ? 

10.  What  kind  of  a  spectrum  is  produced  by  (a}  a  glowing  solid, 
(£)  a  glowing  vapor,  (c)  a  glowing  solid  surrounded  by  a  glowing  vapor? 

n.  What  is  spectrum  analysis  ?  How  is  it  applied  (a)  in  the  labo- 
ratory, and  (b)  by  astronomers  ? 

12.  What  does  spectrum  analysis  show  about  each  element  ?     About 
their  relations  to  each  other  ?     About  their  distribution  ?     About  the 
heavenly  bodies  ? 

13.  Who  perfected  the  spectroscope  and  developed  its  use  ? 

14.  What  recent  use  has  been  made  of  the  spectroscope  in  (a)  chem- 
istry, and  (b)  astronomy  ? 


CHAPTER   XXXI. 
SOME  COMMON  ORGANIC  COMPOUNDS. 

Introduction.  —  In  the  early  days  of  chemistry  it  was 
believed  that  starch,  sugar,  and  other  compounds  obtained 
from  plants  and  animals  were  produced  by  the  influence  of 
some  mysterious  vital  force.  Such  compounds  were  called 
organic,  because  of  their  connection  with  living  things,  i.e. 
with  bodies  having  organs  ;  and  they  were  sharply  dis- 
tinguished from  inorganic  or  mineral  compounds  obtained 
from  the  earth's  crust.  This  distinction  prevailed  until 
Wohler,  in  1828,  prepared  urea  —  a  characteristic  organic 
compound  —  from  inorganic  substances.  Since  then  the 
barrier  between  the  two  classes  of  compounds  has  been 
completely  removed.  We  now  believe  that  compounds  of 
carbon,  whatever  their  source,  are  subject  to  the  laws  that 
govern  all  other  compounds.  The  terms  organic  and  inor- 
ganic are  still  used,  though  they  have  lost  their  original 
narrow  meaning.  Carbon  forms  a  vast  number  of  com- 
pounds which  are  related  to  each  other,  and  which  differ 
markedly  from  most  compounds  of  other  elements.  It  is 
convenient,  therefore,  to  distinguish  these  compounds  by 
the  term  organic  and  to  study  them  under  the  comprehen- 
sive title  of  Organic  Chemistry  or  the  Chemistry  of  Carbon 
Compounds. 

Composition  of  Organic  Compounds.  —  The  number  of 
organic  compounds  is  very  large,  but  they  contain  only  a 
few  elements  —  seldom  more  than  four  or  five.  Hydro- 

405 


406  Descriptive  Chemistry. 

carbons,  as  already  indicated,  contain  carbon  and  hydro- 
gen. Vegetable  substances,  typified  by  starch,  sugar,  and 
fruit  acids,  contain  carbon,  hydrogen,  and  oxygen.  Ani- 
mal substances,  like  hah",  albumen,  gelatine,  and  muscle 
generally  contain  nitrogen  as  well  as  carbon,  hydrogen, 
and  oxygen ;  some  also  contain  sulphur  or  phosphorus. 
Artificial  organic  compounds,  like  dyestuffs,  may  contain 
any  element,  especially  chlorine,  iodine,  and  metals. 

The  number  and  complexity  of  organic  compounds  is 
due  to  several  facts  already  mentioned  in  a  previous 
chapter,  (i)  Atoms  of  carbon  have  power  to  unite  with 
themselves.  (2)  Atoms  of  different  elements  can  be  intro- 
duced into  carbon  compounds.  Sometimes  these  atoms 
are  simply  added,  sometimes  they  replace  other  atoms, 
thus  producing  an  endless  number  of  addition  and  substi- 
tution products.  (3)  The  same  number  of  atoms  may 
arrange  themselves  differently,  thereby  producing  isomeric 
compounds  having  different  properties.  To  these  princi- 
ples, which  should  be  reviewed  until  firmly  grasped,  must 
be  added  another.  (4)  Organic  compounds  contain  radi- 
cals. These  radicals  are  analogous  to  hydroxyl  (OH)  and 
ammonium  (NH4),  and  like  these  radicals  they  exist  only 
in  combination.  They  act  like  single  atoms  and  enter 
unchanged  into  a  number  of  organic  compounds.  The 
radical  C2H5  is  called  ethyl.  It  is  present  in  many 
organic  compounds,  and  its  presence  in  ordinary  alcohol 
gives  rise  to  the  scientific  name,  ethyl  alcohol.  Methyl 
(CH3)  is  another  important  radical,  and  phenyl  (C6H5)  is 
especially  common  in  the  benzene  series  of  organic  com- 
pounds. 

Structure  of  Organic  Compounds.  —  An  extensive  study 
of  the  properties  of  organic  compounds  has  revealed  many 
facts  about  their  constitution,  i.e.  the  structure  of  their 


Some  Common  Organic  Compounds.        407 

molecules.  Little  or  nothing,  of  course,  is  known  about 
the  shape,  size,  etc.,  of  molecules,  but  much  is  known 
about  the  grouping  of  atoms  and  of  radicals  in  the  mole- 
cules. These  facts,  which  are  ascertained  by  experiment 
and  are  often  too  complex  to  be  expressed  briefly,  may  be 
represented  by  suitable  formulas.  The  ordinary  or  empiri- 
cal formula  of  alcohol  is  C2H6O.  But  this  formula  tells 
nothing  about  the  relation  these  atoms  bear  to  each  other, 
nor  whether  all  the  hydrogen  atoms  act  alike.  Experiment 
proves,  however,  that  (i)  one  hydrogen  atom  acts  differ- 
ently from  the  other  five,  and  (2)  one  hydrogen  atom  is 
always  associated  with  the  oxygen  atom  in  chemical 
changes.  Hence,  the  formula  C2H5  .  OH  expresses  more 
fully  these  facts.  Such  a  formula  is  called  a  rational  or 
constitutional  formula.  Sometimes  constitution  is  ex- 
pressed by  a  graphic  formula.  Thus  methane  and  ethane 
have  the  graphic  formulas  — 

H  H     H 

I  I        I 

H— C— H  H— C— C— H 

I  I       I 

H  H     H 

Methane  Ethane 

In  these  diagrams  the  single  lines  represent  a  valence  of 
one  —  nothing  else,  and  the  number  of  lines  connected 
with  each  atom  must  be  equal  to  the  valence  of  the  ele- 
ment in  the  compound.  The  lines  are  sometimes  called 
bonds  or  links,  but  they  are  not  intended  to  represent  at- 
traction or  any  other  force.  Nor  do  they  represent  space 
relations.  In  the  case  of  methane,  they  mean  that  the 
four  hydrogen  atoms  bear  the  same  relation  to  the  single 
carbon  atom.  In  the  case  of  ethane,  they  mean  the  same, 


408  Descriptive  Chemistry. 

but  they  also  indicate  that  the  two  carbon  atoms  are  joined. 
The  graphic  formula  of  ethyl  alcohol  is — 

H     H 

I        I 
H— C  — C— O— H 

I        I 
H     H 

This  is  not  an  arbitrary  arrangement ;  the  facts  mentioned 
above  necessitate  this  general  arrangement.  Additional 
illustrations  of  this  subject  will  be  given,  as  different 
compounds  are  discussed. 

Classification  of  Organic  Compounds.  —  Organic  com- 
pounds are  divided  and  subdivided  into  many  classes 
for  purposes  of  study.  Only  the  most  common  organic 
compounds  can  be  considered  in  this  book.  These  are 
members  of  the  following  groups:  (i)  Hydrocarbons,  (2) 
Alcohols,  (3)  Aldehydes,  (4)  Ethers,  (5)  Acids,  (6)  Ethe- 
real salts,  (7)  Fats,  glycerine,  and  soap,  (8)  Carbohydrates, 
(9)  Benzene  and  its  derivatives.  Some  compounds  are  so 
closely  related  that  they  really  belong  to  several  of  these 
groups,  while  a  few  cannot  strictly  be  put  in  any  of  them. 

HYDROCARBONS. 

Three  of  these  compounds  of  carbon  arid  hydrogen  have 
been  fully  considered  in  Chapter  XV.7  The  chief  facts 
and  fundamental  principles  recorded  there  may  be  profit- 
ably reviewed  at  this  point.  Other  hydrocarbons  will  be 
discussed  under  Benzene  (see  below). 

ALCOHOLS. 

Alcohols  are  compounds  of  carbon,  hydrogen,  and  oxy- 
gen. Ordinary  or  ethyl  alcohol  is  the  best  known  member 


Some  Common  Organic  Compounds.        409 

of  this  group.  It  is  usually  called  simply  alcohol.  There 
are  many  alcohols  analogous  to  ethyl  alcohol,  but  the  only 
other  important  one  is  methyl  alcohol. 

The  alcohols  may  be  regarded  as  hydroxides  of  certain  radicals,  e.g. 
ethyl,  methyl,  propyl,  etc.1  For  example,  ethyl  alcohol  is  ethyl  hydrox- 
ide, and  may  be  considered  as  formed  by  replacing  one  hydrogen  atom 
of  ethane  (C2H(i)  by  one  hydroxyl  group  (OH).  Again,  alcohols  are 
analogous  to  metallic  hydroxides,  in  which  the  metal  is  replaced  by  a 
radicals- 


Ethyl  Hydroxide  Sodium  Hydroxide 

Alcohols  and  metallic  hydroxides  have  some  properties  in  common. 
Thus,  both  form  salts  with  acids.  With  acetic  acid,  sodium  hydroxide 
forms  sodium  acetate,  while  alcohol  forms  ethyl  acetate  (see  Ethereal 
Salts). 

Methyl  Alcohol,  CH3.OH,  is  a  colorless  or  slightly 
yellowish  liquid,  much  like  ordinary  alcohol.  It  boils  at 
about  66°  C,  and  burns  with  a  pale  flame  which  de- 
posits no  soot.  It  intoxicates,  and  if  concentrated  is 
poisonous.  It  mixes  with  water  in  all  proportions.  It  is 
cheaper  than  ethyl  alcohol,  and  is  used  as  a  solvent  for 
fats,  oils,  and  shellac,  and  in  the  manufacture  of  varnishes 
and  dyestuffs.  Methyl  alcohol  is  often  called  wood  alco- 
hol or  wood  spirit,  because  it  is  one  of  the  liquid  products 
obtained  by  the  dry  distillation  of  wood  (see  Charcoal). 

Ethyl  Alcohol,  C2H5.  OH,  is  a  colorless,  volatile  liquid, 
having  a  burning  taste  and  a  pleasant  odor.  It  is  lighter 
than  water,  its  specific  gravity  being  about  0.8.  It  boils 
at  78.3°  C.,  and  does  not  freeze  until  at  —  130.5°  C.  Be- 
cause of  its  very  low  freezing  point,  it  is  used  in  ther- 

1  The  names  of  these  and  similar  radicals  are  derived  from  the  correspond- 
ing hydrocarbon.  Thus,  the  word  methyl  comes  from  methane,  ethyl  from 
ethane,  propyl  from  propane. 


4i o  Descriptive  Chemistry. 

mometers  designed  to  record  temperatures  below  —  40°  C. 
(the  freezing  point  of  mercury),  as  in  Arctic  explorations. 
Its  harmful  effect  on  the  human  system  need  not  be  dis- 
cussed. Alcohol  mixes  with  water  in  all  proportions. 
The  ordinary  commercial  variety  contains  from  50  to  95 
per  cent  of  alcohol.  Pure  or  absolute  alcohol  is  obtained 
by  removing  the  remaining  water  with  lime.  Proof  spirit 
contains  about  50  per  cent  of  alcohol.  Methylated  spirit 
contains  90  per  cent  ethyl  and  10  per  cent  methyl  alcohol; 
it  is  often  used  as  a  cheap  substitute  for  ordinary  alcohol, 
but  it  cannot  be  used  as  a  beverage  on  account  of  the  dis- 
agreeable taste  imparted  by  the  methyl  alcohol.  Alcohol 
is  an  excellent  solvent  for  gums,  oils,  and  resins,  and  is 
therefore  extensively  used  in  the  manufacture  of  varnishes, 
essences,  extracts,  tinctures,  perfumes,  and  medicines.  It 
is  also  used  as  an  antiseptic,  and  as  a  source  of  heat  in 
alcohol  lamps.  Many  organic  compounds,  as  ether  and 
chloroform,  are  prepared  from  alcohol.  Some  vinegar  is 
made  from  alcohol.  In  museums  alcohol  is  used  to  pre- 
serve specimens.  Alcohol  may  be  prepared  from  ethane  (see 
below),  but  it  is  manufactured  by  the  fermentation  of  sugars. 
Fermentation  is  a  general  term  for  the  chemical  changes 
caused  by  ferments.  The  latter  are  usually  minute  living 
bodies,  though  some  inorganic  chemical  sub- 
stances cause  fermentation.  The  process  and 
essential  products  vary  with  the  nature  of  the 
ferment.  The  important  kinds  of  fermenta- 
tion are  alcoholic,  acetic,  and  lactic,  and  the 
respective  products  are  alcohol,  acetic  acid, 
FlG-77-  and  lactic  acid.  Alcoholic  fermentation  is 

Yeast  cells. 

caused  by  ordinary  yeast.  Under  the  micro- 
scope, yeast  has  the  form  of  slimy  yellow  chains  of  small, 
round  cells  (Fig.  77).  When  yeast  is  added  to  a  solution 


Some  Common  Organic  Compounds.        411 

of  glucose,  or  any  other  fermentable  sugar,  the  yeast 
plants  multiply  rapidly.  Air  must  be  admitted,  and  the 
temperature  should  be  2O°-3O°  C.  The  changes  are 
numerous  and  complex,  but  the  main  products  are  alcohol 
and  carbon  dioxide,  thus  — 

C6H1206  2C2H60       +  2C02 

Glucose  Alcohol  Carbon  Dioxide 

The  fermentation  ceases  as  soon  as  the  liquid  contains 
about  14  per  cent  of  alcohol.  The  solution  is  filtered  and 
concentrated  by  distillation,  until  the  distillate  contains  the 
desired  per  cent  of  alcohol.  Commercial  alcohol  is  made 
also  from  potatoes,  grains,  rice,  beet  root,  molasses,  and 
many  other  substances  rich  in  sugar  and  starch.  Ordinary 
or  cane  sugar  must  be  boiled  with  acid  before  it  will 
ferment. 

Wines,  beers,  and  all  alcoholic  liquors  are  prepared  by 
fermentation.  Yeast  is  seldom  added,  however,  because 
the  ferment  which  brings  about  the  change  is  in  the  air, 
upon  fruits  and  vines.  Wines  are  made  from  the  juice  of 
grapes ;  beer  is  made  from  hops  and  malt  (barley  which 
has  sprouted).  Whisky,  gin,  brandy,  rum,  and  cordials 
are  called  distilled  liquors,  and  are  manufactured  by  dis- 
tilling the  liquid  obtained  by  fermenting  grains,  molasses, 
fruit  juices,  and  other  substances  containing  sugar  and 
starches.  Hence,  wine,  beer,  and  similar  liquors  are  essen- 
tially mixtures  of  alcohol  and  water.  They  differ  mainly 
in  their  proportion  of  alcohol.  The  particular  flavor  is  due 
to  small  quantities  of  different  substances  which  are  inten- 
tionally added,  obtained  from  the  raw  materials,  or  formed 
by  special  processes  of  manufacture.  Coloring  matter  is 
usually  added,  but  sometimes  it  is  extracted  from  the  casks 
in  which  the  liquor  is  stored.  Beer  contains  from  3  to  7 


412  Descriptive  Chemistry. 

per  cent  of  alcohol,  wines  from  6  to  20,  rum,  brandy,  and 
whisky  from  40  to  60  or  more  per  cent. 

ALDEHYDES. 

Aldehydes  are  compounds  of  carbon,  hydrogen,  and  oxy- 
gen. They  are  formed  by  the  oxidation  of  alcohols.  The 
two  important  members  of  this  group  are  acetic  aldehyde 
(or  acetaldehyde)  and  formic  aldehyde  (or  formaldehyde). 

Acetic  Aldehyde,  CH3  .  CHO,  is  usually  called  simply  aldehyde.  It 
is  a  colorless,  very  volatile  liquid,  and  has  a  peculiar,  suffocating  odor. 
It  is  a  vigorous  reducing  agent,  and  is  sometimes  used  to  precipitate 
silver,  as  a  thin  coating,  from  silver  solutions.  It  is  converted  by  oxi- 
dizing agents  into  acetic  acid  (hence  its  name,  acetic  aldehyde}.  Alde- 
hyde is  prepared  by  oxidizing  alcohol  with  a  solution  of  potassium  (or 
sodium)  dichromate  and  sulphuric  acid.  When  a  mixture  of  these  three 
substances  is  gently  warmed,  the  characteristic  odor  of  aldehyde  may  be 
detected.  The  oxidation  of  alcohol  consists  simply  in  the  removal  of 
hydrogen,  thus  — 

C2H,.OH     +     O     =     CH3.CHO     +     H2O 
Alcohol  Aldehyde 

The  word  aldehyde  emphasizes  this  fact,  being  a  contraction  of  0/cohol 


When  chlorine  is  used  to  oxidize  alcohol,  part  of  the  hydrogen  is 
replaced  by  chlorine,  and  the  compound  CC13  .  CHO  is  formed.  This 
substance,  called  chloral,  forms  a  hydrate  (CC13  .  CHO  .  H2O),  which  is 
used  to  induce  sleep  and  relieve  pain.  When  chloral  is  treated  with  an 
alkali,  it  is  decomposed  and  chloroform  (CHC13)  is  produced.  The 
latter  is  a  sweet  liquid,  and  is  used  to  produce  insensibility  in  surgical 
operations.  Chloroform  is  usually  made  by  treating  alcohol  with  bleach- 
ing powder.  lodoform  (CHI3),  which  is  analogous  to  chloroform,  is  a 
yellow  solid,  with  a  disagreeable  smell,  and  is  extensively  used  as  a 
dressing  for  wounds.  It  protects  the  wound  from  the  harmful  action  of 
germs. 

Formaldehyde,  H  .  CHO,  is  a  gas,  but  is  used  only  in 
solution.  It  has  a  penetrating  odor.  The  commercial  solu- 


Some  Common  Organic  Compounds.        413 

tion  sold  as  formalin  contains  40  per  cent  of  formaldehyde. 
It  corresponds  to  methane  and  methyl  alcohol,  thus  — 

H  H  H 

I  I  I 

H-C-H  H-C-O-H  C  =  O 

I  I  I 

I  H  H  H 

Methane  Methyl  Alcohol        Formaldehyde 

With  oxygen  it  forms  formic  acid  (hence  its  name,  see 
below).  Large  quantities  of  formaldehyde  are  used  in  the 
manufacture  of  dyestuffs  and  fuming  nitric  acid,  as  a  food 
preservative,  and  a  disinfectant.  When  used  for  the  last 
purpose,  the  solution  is  vaporized  in  a  special  kind  of  lamp, 
and  the  vapors  are  conducted  by  a  small  tube  into  the  room 
to  be  disinfected.  It  is  one  of  the  most  convenient  and 
efficient  of  all  disinfectants,  and  is  very  generally  used. 

I 

ETHERS. 

Ethers  are  compounds  of  carbon,  hydrogen,  and  oxygen. 
They  are  analogous  to  the  metallic  oxides.  They  are 
formed  by  heating  alcohols  with  sulphuric  acid.  Ordinary 
or  ethyl  ether  is  the  best  known  member  of  this  group. 

Ethyl  Ether,  C4H10O,  is  a  colorless,  volatile  liquid,  with 
a  peculiar,  pleasing  taste  and  odor.  It  is  lighter  than 
water,  its  specific  gravity  being  about  0.74.  It  boils  at 
35°  C,  and  the  vapor  is  very  inflammable.  The  liquid 
should  never  be  brought  near  a  flame.  It  is  somewhat 
soluble  in  water,  and  it  also  dissolves  water  to  a  slight 
extent.  It  mixes  with  alcohol  in  all  proportions.  It  is  a 
good  solvent  for  waxes,  fats,  oils,  and  other  organic  com- 
pounds. Its  chief  use  is  as  an  anaesthetic,  i.e.  to  render  one 
insensible  to  pain  in  surgical  operations. 


414  Descriptive  Chemistry. 

Ether  is  manufactured  by  distilling  a  mixture  of  ethyl  alcohol  and 
sulphuric  acid  in  the  proper  proportions.  Hence,  the  names,  —  ethyl 
or  sulphuric  ether.  Ethylsulphuric  acid  is  first  produced,  thus  — 

C2H5.OH     +  H2S04  HC2H5SO4         +     H2O 

Alcohol  Sulphuric  Acid  Ethylsulphuric  Acid 

When  more  alcohol  and  the  ethylsulphuric  acid  are  heated  together, 
ether  is  formed,  and  sulphuric  acid  is  reproduced,  thus,  — 

HC2H5SO4  +  C2H5 .  OH  =  (C2H5)2O  +  H2SO4 

Ether 

The  process  is  thus  continuous,  a  small  quantity  of  sulphuric  acid  serv- 
ing to  transform  a  large  quantity  of  alcohol  into  ether.  Ethyl  ether  is 
ethyl  oxide,  (C2H5)2O  or  C2H5  .  O  .  C2H5. 

ACIDS. 

Organic  Acids  are  compounds  of  carbon,  hydrogen,  and 
oxygen.  It  is  a  large  class  of  compounds  divided  into 
several  series,  one  of  the  most  important  of  which  is  the 
acetic  or  fatty  series.  Its  best  known  member  is  acetic 
acid ;  several  of  the  higher  members  occur  in  fats  and  oils. 

These  acids  are  closely  related  to  hydrocarbons,  alcohols,  and  alde- 
hydes, as  may  be  seen  by  the  following  formulas  :  — 

H  H 

I  I 

H-C-H  H-C-(OH) 

I  I 

H-C-H  H-C-H 

I  I 

H  H 

Ethane  Ethyl  Alcohol  Acetic  Aldehyde      Acetic  Acid 

It  is  thus  possible  to  pass  from  a  hydrocarbon  through  a  correspond- 
ing alcohol  and  aldehyde  to  an  acid. 

The  characteristic  group  of  atoms  in  organic  acids  is  COOH  (or 
O  =  C  -  O  -  H),  and  is  called  carboxyl. 


H 

1 

C  =  O 

O=C-(OH) 

1 

1 

H-C-H 

H-C-H 

1 

1 

H 

H 

Some  Common  Organic  Compounds.      ,415 

Acetic  Acid,  C2H4O2  or  CH3.  COOH.— This  is  the 
most  common  organic  acid.  It  is  manufactured  on  a  large 
scale  by  the  dry  distillation  of  wood.  The  dark  red 
watery  distillate,  which  is  called  pyroligneous  acid,  con- 
tains about  10  per  cent  of  acetic  acid  besides  a  small  per 
cent  of  methyl  alcohol  and  many  other  organic  compounds. 
This  distillate  is  neutralized  with  lime  or  sodium  carbonate, 
and  the  acetate  formed  is  then  decomposed  and  distilled 
with  hydrochloric  or  sulphuric  acid.  The  acetic  acid  which 
condenses  in  the  receiver  may  be  further  purified  by  dis- 
tilling it  with  potassium  dichromate  and  then  filtering 
through  charcoal.  Sometimes  the  pyroligneous  acid  is 
distilled  without  neutralizing ;  the  distillate  is  then  dilute, 
impure  acetic  acid,  known  as  wood  vinegar.  If  sodium 
acetate,  prepared  as  described  above,  is  fused  and  then 
distilled  with  concentrated  sulphuric  acid,  the  product  is 
a  very  concentrated  acetic  acid.  It  is  called  glacial  acetic 
acid,  because  at  about  1 7°  C.  it  becomes  an  icelike  solid. 

Commercial  acetic  acid  is  a  water  solution  containing 
about  30  per  cent  of  pure  acetic  acid.  It  is  a  colorless 
liquid,  having  a  pleasant  odor  and  a  sharp  taste.  It  is 
slightly  heavier  than  water.  It  mixes  with  water  and  alco- 
hol in  all  proportions,  and  like  alcohol  is  an  excellent 
solvent  for  many  organic  substances.  Recently,  it  has 
begun  to  replace  alcohol  as  a  solvent  for  many  drugs. 

Acetic  acid  is  used  to  prepare  acetates,  dyestuffs,  and 
other  organic  compounds,  medicines,  white  lead,  and  in 
the  manufacture  of  vinegar. 

Vinegar  is  dilute,  impure  acetic  acid.  It  is  prepared  by 
oxidizing  dilute  alcohol,  the  essential  change  being  repre- 
sented thus  — 

C2H60     4-     02     =     C2H402     +     H20 
Alcohol  Oxygen          Acetic  Acid  Water 


416 


Descriptive  Chemistry. 


The  transformation  is  accomplished  by  fermentation. 
Two  processes  are  used,  (i)  When  beer,  weak  wines,  or 
cider  are  exposed  to  the  air,  they  slowly  become  sour, 
owing  to  the  conversion  of  alcohol  into  acetic  acid.  The 
change  is  caused  by  the  presence  and  activity  of  a  ferment, 
known  as  mycoderma  aceti,  or  "  mother  of  vinegar."  Strong 
wines  and  pure  dilute  alcohol  do  not  become  sour,  because 
the  ferment  cannot  live  in  such  liquids.  (2)  In  the  "quick 
vinegar  process,"  impure  dilute  alcohol  is  oxidized  by  ex- 
posing it  to  an  excess  of  air.  The  operation  is  conducted 

in  tall  vats  or  casks  filled  with 
beechwood  shavings  soaked 
in  strong  vinegar  (Fig.  78). 
Holes  at  the  bottom  and  top 
allow  air  to  enter  and  escape 
freely.  The  alcoholic  solu- 
tion is  introduced  at  the  top, 
trickles  through  the  shavings, 
and  collects  at  the  bottom. 
In  its  passage  it  comes  in 
contact  with  the  ferment  and 
oxygen,  and  is  partially  con- 
verted into  vinegar.  The 
operation  is  repeated  until 
Thus  prepared,  the  vinegar 

lacks  the  flavor,  odor,  and  color  of  cider  vinegar,  but  these 
deficiencies  are  often  artificially  supplied. 

Vinegar  is  used  chiefly  as  a  condiment  for  the  table  and 
in  making  pickles  and  similar  relishes. 

The  constitution  of  acetic  acid  has  been  shown  to  correspond  to  the 
formula  CH3  .  COOH.     Its  metallic  salts  are  formed  by  substituting  a 
metallic  atom  (or  group)  for  the  hydrogen  of  the  group  COOH. 
radical  CH3  remains  unchanged.     (See  page  170.) 


FIG.  78.  —  Apparatus  for  the  prep 
aration  of  vinegar  from  impure,  dilute 
alcohol. 

the    change   is   complete. 


The 


Some  Common  Organic  Compounds.        417 

Acetates.  —  Acetic  acid  is  a  monobasic  acid,  and  forms 
a  series  of  salts  —  the  acetates.  They  are  prepared  like 
other  salts  by  the  interaction  of  the  acid  and  carbonates, 
hydroxides,  metals,  etc.  The  metallic  acetates  are  usually 
crystallized  solids,  which  readily  yield  acetic  acid  when 
treated  with  sulphuric  or  a  similar  acid.  Most  of  them 
contain  water  of  crystallization,  and  most  are  poisonous. 

Several  acetates  have  useful  applications.  Sodium  acetate, 
NaC2H3O2 .  3  H2O,  is  a  white  crystallized  solid,  used  in  preparing 
pure  acetic  acid,  and  in  the  manufacture  of  dyestuffs.  Lead  acetate, 
Pb(C2HsO2)2,  is  a  white  crystallized  solid,  used,  in  dyeing  and  in  mak- 
ing a  yellow  pigment.  Its  sweet  taste  led  to  the  common  name  of 
"sugar  of  lead.11  Aluminium  acetate,  A1(C2H3O2)3,  is  not  known  in 
the  pure  state,  but  an  impure  solution,  known  as  "  red  liquor,"  is  exten- 
sively used  in  dyeing  and  calico  printing.  Iron  acetates  are  sold  in 
solution  as  a  complex  black  liquid,  known  as  "iron  liquor,"  which  is 
used  in  dyeing  black  silks  and  cottons,  and  in  calico  printing  (see 
Mordants).  A  complex  copper  acetate,  2  Cu(C2H3O2)2  +  CuO,  called 
verdigris,  is  used  in  making  blue  paint.  Another  complex  acetate  of 
copper  and  arsenic  is  Paris  green ;  it  is  used  to  kill  potato  bugs  and 
other  insects  which  injure  vegetation. 

A  few  other  acids  in  this  series  are  interesting.  Butyric  acid 
C4H8O9,  is  the  acid  which  gives  the  disagreeable  odor  to  rancid  butter. 
Stearic  acid.  C18H3(JO2,  and  Palmitic  acid,  C16Ho2O2,  are  found  as 
compounds  in  beef  suet,  mutton  fat,  butter,  and  other  fats.  Palmitic 
acid  is  also  one  of  the  essential  compounds  found  in  palm  oil.  These 
two  acids  are  white  solids,  and  are  used  to  make  stearin  candles  (see 
Fats,  below). 

Other  Organic  Acids  which  are  important  are  oxalic, 
lactic,  malic,  tartaric,  and  citric. 

Oxalic  Acid  occurs  as  a  salt  in  rhubarb  and  sorrel.  It 
is  manufactured  on  a  large  scale  by  heating  sawdust  with 
potassium  hydroxide,  and  treating  the  residue  first  with 
lime  and  then  with  sulphuric  acid.  Oxalic  acid  is  a  white 
solid,  very  soluble  in  water,  from  which  it  crystallizes  with 


4i  8  Descriptive  Chemistry. 

two  molecules  of  water  of  crystallization  (C2H2O4  .  2.  H2O). 
It  is  very  poisonous.  It  is  dibasic  and  forms  several  use- 
ful salts.  The  acid  and  some  of  its  salts  decompose  iron 
rust  and  inks  containing  iron,  and  are  often  used  to  remove 
such  stains  from  cloth.  The  acid  and  its  salts  are  also 
used  in  dyeing,  calico  printing,  photography,  in  making 
dyestuffs,  and  as  an  ingredient  of  mixtures  for  cleaning 
brass  and  copper. 

Lactic  Acid,  C3H6O3,  occurs  in  sour  milk,  being  one 
product  of  the  fermentation  of  the  milk  sugar.  It  is  a 
thick,  sour  liquid,  and  is  easily  decomposed  by  heat. 
When  sour  milk  is  used  in  cooking,  the  "  baking  soda  " 
and  lactic  acid  interact,  producing  soluble  sodium  lactate 
and  carbon  dioxide  gas.  Lactic  acid  and  its  salts  are  used 
as  medicines,  in  beverages,  and  as  a  substitute  for  more 
expensive  acids  in  dyeing  and  calico  printing. 

Malic  acid,  C4H6O5,  is  found  free  and  as  salts  in  apples,  pears,  cur- 
rants, gooseberries,  rhubarb,  grapes,  and  berries  of  the  mountain  ash 
tree.  It  is  a  white,  crystalline  solid. 

Tartaric  Acid,  C4H6O6,  occurs  as  the  potassium  salt  in 
grapes  and  other  fruits.  During  the  fermentation  of  grape 
juice,  impure  acid  potassium  tartrate  is  deposited  in  the 
casks.  From  this  argol  or  crude  tartar  the  acid  itself 
is  prepared  by  treating  the  raw  product  successively  with 
chalk  and  sulphuric  acid.  Tartaric  acid  is  a  white  crystal- 
lized solid,  soluble  in  water  and  alcohol.  It  is  used  in  dye- 
ing, and  as  one  ingredient  of  Seidlitz  powders.  In  these 
and  similar  powders  it  serves  to  decompose  the  other  in- 
gredient which  is  a  carbonate  (see  Sodium  Bicarbonate). 

Tartaric  acid  is  dibasic  and  forms  two  classes  of  salts.  Purified 
acid  potassium  tartrate  obtained  from  argol  is  commonly  known  as 
cream  of  tartar.  It  is  extensively  used  in  the  manufacture  of  baking 
powders.  These,  as  a  rule,  are  essentially  mixtures  of  cream  of  tartar 


Some  Common   Organic  Compounds.        419 

and  sodium  bicarbonate,  HNaCO3.  When  moistened  by  dough,  the 
baking  powder  dissolves,  the  two  ingredients  interact  and  liberate  car- 
bon dioxide  as  the  main  product.  This  gas  bubbles  slowly  through 
the  dough,  thereby  puffing  it  up  and  making  it  porous  (see  Sodium 
Bicarbonate).  Tartar  emetic  is  a  tartrate  of  potassium  and  antimony. 
It  is  used  as  a  medicine  and  to  some  extent  in  dyeing. 

Citric  Acid,  C(;H8O7,  occurs  abundantly  in  lemons  and  oranges,  and 
in  small  quantities  in  currants,  gooseberries,  and  raspberries.  It  is  a 
white,  crystallized  solid,  very  soluble  in  water.  The  taste  is  sour,  but 
pleasant.  The  acid  and  its  magnesium  salt  are  used  as  medicines.  The 
acid  itself  is  used  in  calico  printing.  Citric  acid  is  tribasic. 

ETHEREAL    SALTS. 

Ethereal  Salts  or  Esters  are  compounds  of  carbon,  hy- 
drogen, and  oxygen  closely  related  to  alcohols  and  organic 
acids.  Thus,  when  ethyl  alcohol,  acetic  acid,  and  concen- 
trated sulphuric  acid  are  mixed  and  warmed,  ethyl  acetate 
is  formed.  The  essential  change  is  represented  thus  — ^ 

C2H5.OH   +CH3.COOH   =   CH3.COOC2H5+   H2O 

Ethyl  Alcohol  Acetic  Acid  Ethyl  Acetate  Water 

The  sulphuric  acid  serves  to  absorb  the  water.  Ethyl 
acetate  has  a  pleasant,  fruitlike  odor,  and  its  formation  in 
this  way  is  a  simple  test  for  alcohol  or  acetic  acid.  Ethyl 
acetate  is  analogous  to  sodium  acetate,  i.e.  the  organic  salt 
contains  the  radical  ethyl  while  the  metallic  salt  con- 
tains sodium.  The  fatty  acids,  as  well  as  those  of  other 
series,  form  many  ethereal  salts  of  special  interest.  Some 
occur  naturally  in  fruits  and  flowers,  and  in  many  cases 
give  the  flavor  and  fragrance.  Others  are  prepared  artifi- 
cially and  used  as  the  basis  of  cheap  flavoring  extracts, 
perfumery,  and  beverages.  Ethyl  butyrate  has  the  taste 
and  fragrance  of  pineapples,  amyl  acetate  of  bananas, 
amyl  valerate  of  apples. 


420  Descriptive  Chemistry. 

FATS,  GLYCERINE,  AND  SOAP. 

General  Relations.  —  Natural  fats  and  oils  are  essentially 
mixtures  of  stearin,  palmitin,  and  olein.  Beef  and  mutton  fat 
are  chiefly  stearin,  lard  is  mainly  palmitin  and  olein  ;  while 
oils,  such  as  olive  oil,  are  largely  olein.  Stearin  and  pal- 
mitin are  solids  at  the  ordinary  temperature,  but  olein  is  a 
liquid.  These  three  compounds  —  stearin,  palmitin,  and 
olein  —  are  ethereal  salts  of  their  corresponding  acids  and 
the  alcohol,  glycerine.  They  are  analogous  to  ethyl  acetate. 
The  radical  of  glycerine  is  glyceryl,  C3H5.  Thus,  stearin 
is  glyceryl  stearate,  palmitin  is  glyceryl  palmitate,  and 
olein  is  glyceryl  oleate.  Natural  fats  and  oils,  therefore, 
are  mixtures  of  these  and  similar  ethereal  salts.  Fats  are 
sometimes  called  glycerides.  Glycerine  is  a  triacid  alcohol 
containing  three  hydroxyl  (OH)  groups.  Like  ordinary 
alcohol,  it  interacts  with  the  fatty  acids  and  forms  ethereal 
salts.  The  latter,  as  we  have  just  learned,  are  the  fats. 
Now  when  fats  are  heated  with  very  hot  steam  or  with  sul- 
phuric acid,  the  fats  themselves  are  changed  into  glycerine 
and  the  corresponding  acids.  Thus,  with  stearin,  the 
change  is  — 

(C17H35  .  C02)3C3H5  +  3  H20   =   C3H5(OH),   +   3  CirH,, .  COOH 

Stearm  Glycerine  Stearic  Acid 

But  if  fats  are  boiled  with  sodium  hydroxide  or  a  simi- 
lar alkali,  glycerine  and  an  alkaline  salt  of  the  correspond- 
ing acid  are  formed.  Soap  is  a  mixture  of  such  alkaline 
salts.  In  a  few  words,  the  general  relations  are  these: 
(i)  fats  are  ethereal  salts.  (2)  Treated  with  steam  or  acid, 
fats  form  glycerine  and  fatty  acids.  (3)  Treated  with  alka- 
lies, fats  form  glycerine  and  soap. 

Natural  Fats  and  Oils  are  often  complicated  mixtures. 
The  solid  fats,  as  already  stated,  are  rich  in  stearin  and 


Some  Common  Organic  Compounds.        421 

palmitin.  Tallow  is  chiefly  stearin,  but  human  fat  and 
palm  oil  are  largely  palmitin.  The  soft  and  liquid  fats  and 
oils  contain  considerable  olein,  as  a  rule.  The  proportion 
of  olein  determines  the  consistency  of  the  fats  and  oils. 
Thus,  Olive  oil  contains  about  72  per  cent  of  olein  (and  a 
similar  fat)  and  28  per  cent  of  stearin  and  palmitin.  The 
specific  character  of  many  fats  and  oils  is  due  mainly  to 
the  presence  of  a  small  proportion  of  certain  fats.  These 
fats  correspond  to  uncommon  acids  in  the  fatty,  oleic,  and 
other  series.  Butter,  for  example,  consists  mainly  of  the 
fats  corresponding  to  the  following  acids  :  palmitic,  stearic, 
oleic,  butyric,  capric,  and  caproic.  The  last  three  with 
traces  of  other  substances  give  butter  its  pleasant  flavor. 
Oleomargarine  and  other  substitutes  for  butter  resemble 
real  butter  very  closely  in  composition.  Artificial  butter, 
however,  lacks  the  flavor  of  the  real  butter,  but  it  is  "  prob- 
ably just  as  nutritious,  although  perhaps  not  quite  so  easily 
digested."  The  lack  of  flavor  noticed  in  artificial  butter  is 
due  to  the  absence  of  the  fats  corresponding  to  the  acids 
of  low  molecular  weight.  Cottolene  is  a  mixture  of  beef 
fat  and  cotton-seed  oil ;  it  is  used  as  a  substitute  for  lard. 

Glycerine  (C3H8O3  or  C3H5.(OH)3)  is  a  thick,  sweet 
liquid.  It  mixes  readily  with  water  and  with  alcohol  in  all 
proportions,  and  absorbs  moisture  from  the  air.  Heated 
in  the  air,  it  decomposes  and  gives  off  irritating  gases,  like 
those  produced  by  burning  fat. 

Glycerine  is  used  to  make  nitroglycerine  (see  below), 
toilet  soaps,  printers'  ink  rolls ;  it  is  also  used  as  a  solvent, 
a  lubricator,  a  preservative  for  tobacco  and  certain  foods, 
a  sweetening  substance  in  certain  liquors,  preserves,  and 
candy ;  as  a  cosmetic ;  and,  owing  to  its  non-volatile  and 
non-drying  properties,  it  is  used  as  an  ingredient  of  inks 
and  oils. 


422  Descriptive  Chemistry. 

Glycerine  is  a  by-product  in  the  manufacture  of  soap,  or  it  is  made 
directly  by  decomposing  fats  with  steam  under  pressure  or  with  lime. 
Ail  these  methods  involve  the  chemical  change  described  above,  viz. 
the  decomposition  of  an  ethereal  salt  (the  fat)  into  the  corresponding 
alcohol  (glycerine)  and  a  mixture  of  fatty  acids.  By  skillful  treatment 
the  glycerine  is  freed  from  water  and  impurities.  The  mixture  of  fatty 
acids  is  made  into  the  so-called  "stearin"  candles. 

As  already  stated,  glycerine  is  an  alcohol,  and  for  this  reason  it  is 
often  called  glycerol.  When  treated  with  a  mixture  of  concentrated 
nitric  and  sulphuric  acids,  it  forms  an  ethereal  salt  commonly  known  as 
nitroglycerine  (C3H3(ONO2)3).  This  is  a  yellow,  heavy,  oily  liquid. 
It  is  the  well-known  explosive,  and  is  also  an  ingredient  of  some  other 
explosives.  When  kindled  by  a  flame,  it  burns  without  explosion ;  but 
if  struck  by  a  hammer  or  heated  suddenly  by  a  percussion  cap,  it  ex- 
plodes violently.  Nitroglycerine  is  used  in  blasting  ;  but  since  it  is  dan- 
gerous to  handle  and  transport,  it  is  usually  mixed  with  some  porous 
substance,  such  as  infusorial  earth,  fine  sand,  or  even  sawdust.  In  this 
form  it  is  called  dynamite. 

Soap,  as  already  stated,  is  a  mixture  of  alkaline  salts  of 
organic  acids,  mainly  stearic  and  palmitic  acids.  Soap  is 
made  by  boiling  fats  with  sodium  hydroxide  or  potassium 
hydroxide.  This  process  is  called  saponification.  Sodium 
hydroxide  produces  hard  soap,  consisting  chiefly  of  sodium 
palmitate,  sodium  stearate,  and  sodium  oleate.  Potassium 
hydroxide  produces  soft  soap,  which  is  mainly  the  corre- 
sponding potassium  salts.  The  chemical  change,  as  already 
stated,  consists  in  tr  e  transformation  of  an  ethereal  salt 
(fat)  into  glycerine  and  an  alkaline  salt.  In  the  case  of 
pure  stearin  (glyceryl  stearate)  the  change  may  be  repre- 
sented thus  — 

C3H5(C17H35 .  C02)3  +  3NaOH   -  3  C17H35 .  CO2Na   +    C3H5(OH)3 
Stearin  Sodium  Sodium  Glycerine 

Hydroxide  Stearate 

The  fats  used  in  soap  making  vary  with  the  soap.  Tal- 
low, lard,  palm  oil,  and  cocoanut  oil  make  white  soaps. 


Some  Common  Organic  Compounds.        423 

Bone  grease  or  house  grease,  together  with  tallow,  palm 
oil,  cotton-seed  oil,  and  rosin,  make  yellow  soaps.  Olive 
oil  is  used  for  making  castile  soap. 

In  the^cold  process  the  calculated  amounts  of  alkali  and  fat  are  allowed 
to  interact,  first  in  a  large  tank  and  then  in  a  box  called  a  "  frame."  By 
this  process  the  glycerine  and  excess  of  alkali  are  left  in  the  soap.  Most 
soaps  are  made  by  the  boiling  process.  The  fat  and  alkali  are  boiled 
in  a  huge  kettle.  This  operation  produces  a  thick,  frothy  mixture  of 
soap,  glycerine,  and  alkali.  At  the  proper  time,  salt  is  added,  thereby 
causing  the  soap  to  separate  and  rise  to  the  top.  The  liquid  beneath  is 
drawn  off,  and  from  it  glycerine  is  extracted.  The  soap  is  often  boiled 
again  with  rosin  or  cocoanut  oil ;  then  purified  by  washing,  mixed,  if 
desired,  with  perfume,  coloring  matter,  or  some  filling  material  (such  as 
sodium  silicate,  sand,  borax),  cooled  in  "frames,"  cut,  and  dried.  Most  • 
soaps  contain  water.  This  really  assists  their  cleansing  action.  The 
latter  is  believed  to  be  due  to  the  free  alkali  formed  by  the  decomposi- 
tion of  the  soap  when  dissolved. 

CARBOHYDRATES. 

Carbohydrates  are  compounds  of  carbon,  hydrogen,  and 
oxygen.  This  is  a  large  group,  and  the  most  important 
members  are  the  sugars,  starches,  and  cellulose. 

The  term  carbohydrate  is  applied  to  these  compounds  because  they 
contain  hydrogen  and  oxygen  in  the  proportion  to  form  water.  They 
were  once  regarded  as  hydrates  of  carbon,  or  carbon  hydrates  —  a  view 
which  is  incorrect  and  misleading. 

Sugars.  —  The  popular  term  sugar  means  almost  any 
sweet  substance  found  in  fruits,  nuts,  vegetables,  sap  of 
trees,  etc.,  though  it  is  usually  restricted  to  the  ordinary 
white  sugar  obtained  from  sugar  cane  and  sugar  beet. 
Chemically,  there  are  many  sugars,  each  having  a  defi- 
nite constitution.  The  most  important  is  ordinary  sugar, 
which  is  also  called  cane  sugar,  sucrose,  and  saccharose. 
Another  important  sugar  is  glucose. 


424  Descriptive  Chemistry. 

Cane  Sugar,  C12H22On,  is  widely  distributed  in  nature, 
being  found  in  the  sugar  cane,  sugar  beet,  sugar  maple, 
Indian  corn,  sorghum,  most  sweet  fruits,  many  nuts,  blos- 
soms of  flowers,  and  honey.  The  main  source  of  cane 
sugar  is  the  sugar  cane  and  sugar  beet. 

Saccharose,  or  ordinary  sugar,  is  a  white,  crystallized 
solid.  Rock  candy  is  highly  crystallized  sugar.  It  is  solu- 
ble in  water,  but  only  sparingly  soluble  in  alcohol.  Heated 
to  160°  C,  sugar  melts,  and  on  cooling  forms  a  pale  yellow 
colored  mass,  called  barley  sugar.  Heated  to  about  200°  C., 
it  is  changed  into  water  and  a  brown  mass,  called  caramel, 
which  is  used  to  color  liquors,  soups,  etc.  If  sugar  is 
heated  with  sulphuric  acid,  it  is  changed  into  a  black  mass, 
which  is  mainly  carbon ;  several  gases  are  also  produced, 
such  as  steam,  carbon  dioxide,  and  sulphur  dioxide.  Cane 
sugar  does  not  ferment. 

The  manufacture  of  Cane  Sugar  from  sugar  cane  and  sugar  beets 
involves  two  main  operations:  (i)  the  preparation  of  raw  sugar  and 
(2)  its  purification  or  refining,  (i)  In  the  preparation  of  raw  sugar 
from  sugar  cane  the  juice  is  extracted  from  the  cane  by  crushing  the 
latter  between  heavy  iron  rollers.  The  liquid  is  then  clarified  as  soon 
as  possible  by  boiling  it  with  a  little  lime,  removing  the  scum  which 
contains  much  of  the  impurity,  and  finally  filtering  the  liquid  through 
bags  or  a  filter  press.  The  purified  juice  is  next  evaporated  until  the 
cane  sugar  begins  to  crystallize  from  the  cooled  liquid.  Formerly  the 
evaporation  was  accomplished  in  an  open  pan,  and  is  now  in  some 
localities,  but  usually  a  vacuum  kettle  is  used.  The  crystals  are  next 
separated  from  the  liquid  by  allowing  the  latter  to  drip  out,  or  more 
commonly  by  whirling  it  out  in  a  centrifugal  machine.  The  solid 
product  is  called  muscovado,  raw  or  brown  sugar.  The  thick  liquid 
is  the  familiar  molasses.  There  are  several  grades  of  each  product. 
The  preparation  of  raw  sugar  from  sugar  beets  resembles  the  method 
used  for  sugar  cane.  The  washed  beets  are  reduced  to  a  pulp,  or  cut 
into  slices,  and  then  treated  with  water.  The  sugar  dissolves  in  the 
water.  The  solution  is  clarified,  evaporated,  and  separated  by  pro- 
cesses much  like  those  applied  to  cane-sugar  solutions.  The  raw  sugar 


Some  Common  Organic  Compounds.        425 

can  scarcely  be  distinguished  from  cane  sugar.  The  molasses  is  unfit 
for  table  use,  though  considerable  sugar  is  extracted  from  it  by  means 
of  strontium  hydroxide  (see  Strontium  Hydroxide).  (2)  Raw  sugar  is 
usually  dark  colored,  and  must  be  refined  before  it  is  suitable  for  most 
uses.  The  refining  of  sugar  consists  in  (a)  purification,  and  (<£)  recrys- 
tallization.  («)  The  raw  sugar  is  purified  by  first  dissolving  it  in  huge 
tanks.  Air  is  blown  in  to  agitate  the  heated  solution,  blood  and  other 
substances  are  often  added  to  entangle  the  impurities,  and  lime  is  also 
added  to  precipitate  and  gather  the  impurities  into  a  scum  or  clot. 
The  colored  liquid  is  next  filtered,  first  through  cloth  bags  and  then 
through  animal  charcoal,  from  which  it  drips  as  a  perfectly  clear  liquid, 
(b}  The  filtered  sirup  is  now  evaporated  in  a  large  vacuum  kettle. 
When  a  sample  shows  that  the  evaporation  has  reached  the  proper 
point,  the  liquid  is  run  into  tanks  to  crystallize.  The  crystals  of  sugar 
are  separated  from  the  sirup  by  centrifugal  machines.  The  latter  is 
boiled  again  or  sold  as  sirup  for  the  table.  The  crystals  are  dried  in  a 
heated  tube  called  a  granulator,  so  that  each  grain  will  be  separate. 
Hence  the  name  granulated  sugar.  The  grains  are  sifted  and  packed 
in  barrels  for  the  market. 

Lactose,  or  sugar  of  milk,  has  the  same  formula  as  cane  sugar,  but 
its  constitution  and  properties  differ.  It  is  obtained  from  milk.  Its 
crystals  are  white,  hard,  gritty,  less  sweet  than  cane  sugar ;  they  con- 
tain one  molecule  of  water  of  crystallization.  Sugar  of  milk  is  used  in 
making  homeopathic  pills  and  certain  kinds  of  foods  for  infants. 

Glucose  is  the  name  of  a  sugar  and  of  a  commercial 
mixture  of  glucose  and  several  related  substances.  Glu- 
cose (dextrose  or  grape  sugar,  C6H12O6)  is  found  in  many 
sweet  fruits,  especially  in  grapes.  Old  raisins  are  some- 
times coated  with  this  sugar.  It  is  often  associated  with 
levulose  (fructose  or  fruit  sugar)  —  an  is6meric  compound 
(C6H12O6).  The  two  sugars  are  found,  for  example,  in 
honey  and  in  parts  of  some  plants.  Both  sugars  are 
formed  from  cane  sugar  by  boiling  it  with  a  dilute  acid. 
The  chemical  change  may  be  represented  thus  — 

C12H22On     +    H2O    =    C6Hi2O6  +   C6H12O6 
Cane  Sugar  Glucose  Fructose 


426  Descriptive  Chemistry. 

Both  glucose  and  fructose  ferment,  forming  alcohol  and 
carbon  dioxide  (see  Alcohol). 

The  commercial  mixture  called  "glucose"  is  prepared  on  a  large 
scale  by  boiling  starch  with  a  dilute  acid,  usually  sulphuric  acid.  The 
consistency  and  composition  of  the  product  vary  with  the  details  of 
manufacture.  The  liquid  products  are  called  "  glucose  "  or  "mixing 
sirup,"  while  the  solid  product  is  known  as  "  grape  sugar "  or  "  dex- 
trose." All  contain  more  or  less  glucose  and  are  about  three  fifths  as 
sweet  as  sugar.  But  since  they  dissolve  in  water,  and  are  cheaper  than 
cane  sugar,  they  are  used  extensively  in  the  manufacture  of  candy,  jelly, 
table  sirups,  etc.  They  are  also  added  to  wines  and  liquors,  certain 
medicines,  and  many  thick  liquids  in  which  their  presence  is  harmless. 
In  alkaline  solutions,  glucose  is  a  strong  reducing  agent,  and  is  used  as 
such  in  dyeing  with  indigo.  It  also  reduces  an  alkaline  mixture  of  cop- 
per sulphate,  known  as  Fehling's  solution.  When  this  solution  is 
boiled  with  glucose,  a  reddish  copper  compound  (cuprous  oxide)  is 
formed.  The  presence  of  sugar  in  solution  is  often  shown  in  this  way. 

Starch  is  widely  distributed  in  the  vegetable  kingdom. 
It  is  found  in  wheat,  corn,  and  all  other  grains,  in  pota- 
toes, beans,  peas,  and  similar  vegetables,  and  in  large 
quantities  in  rice,  sago,  tapioca,  and  nuts.  Many  parts  of 
plants  contain  starch,  for  example,  the  stalk,  stem,  leaves, 
root,  seed,  and  fruit.  The  food  value  of  vegetables  de- 
pends largely  upon  the  starch  they  contain. 


FIG.  79.  —  Starch  grains  (magnified) — wheat  (left),  rice  (center),  corn  (right). 

Starch  is  a  white  powder,  as  usually  seen.     But  under 
the  microscope  it  is  found  to  consist  of  a  mass  of  oval 


Some  Common  Organic  Compounds.        427 

grains,  varying  somewhat  with  the  source  (Fig.  79). 
Starch  is  only  very  slightly  soluble  in  water.  But  if 
heated  with  water,  the  grains  swell  and  burst,  partially 
dissolve,  and  form  a  solution  which,  when  cold,  becomes 
the  familiar  starch  paste.  Starch  in  solution  is  turned 
blue  by  iodine,  and  its  presence  in  many  vegetables  and 
foods  may  be  readily  shown  by  grinding  the  substance  in 
a  mortar  with  warm  water  and  adding  a  drop  of  iodine 
solution. 

Starch  is  prepared  on  a  large  scale  chiefly  from  corn  and  potatoes. 
The  operation  is  mainly  mechanical,  and  consists  in  separating  the 
starch  from  the  fatty,  nitrogenous,  and  mineral  matters  in  the  raw 
product.  Immense  quantities  are  consumed  as  food,  in  laundries,  in 
finishing  cloth  and  paper,  in  making  glucose,  and  as  a  paste. 

The  composition  of  starch,  according  to  some  authorities,  corre- 
sponds to  the  formula  C6H10O,,  but  its  formula  is  still  being  investigated. 

Dextrin  is  a  sticky  solid  formed  from  starch  by  heating 
it  to  2OO°-25O°  C.  or  by  treating  it  with  dilute  acids.  It 
is  soluble  in  water  and  forms  a  sticky  solution.  Commer- 
cial dextrin  or  British  gum  is  a  mixture  of  dextrin  and 
similar  compounds.  Mucilage  contains  dextrin.  Large 
quantities  are  used  as  the  gum  for  the  backs  of  postage 
stamps,  and  for  sticking  the  colors  to  the  cloth  in  calico 
printing. 

Dextrin  is  sometimes  regarded  as  an  intermediate  product  between 
starch  and  dextrose.  Its  composition,  according  to  some  authorities, 
corresponds  to  the  formula  C12H20O10,  but  the  statement  made  about  the 
composition  of  starch  also  applies  to  dextrin. 

Bread.  —  Wheat  flour  contains  about  70  per  cent  of  starch.  The  re- 
mainder is  chiefly  water  and  gluten  in  nearly  equal  proportions,  though 
small  quantities  of  mineral  matter,  dextrin,  and  other  fermentable  sub- 
stances are  present.  In  making  bread,  flour,  milk  or  water,  and  a  little 
yeast  are  thoroughly  mixed  into  dough,  which  is  put  in  a  warm  place  to 
rise,  Fermentation  begins  at  once.  The  yeast  changes  the  ferment- 


428  Descriptive  Chemistry. 

able  substances  into  alcohol  and  carbon  dioxide.  The  gases,  in  trying 
to  escape,  puff  up  the  dough,  which  literally  rises  and  becomes  light  and 
porous.  When  the  dough  is  baked,  the  heat  kills  the  yeast,  and  fer- 
mentation stops  ;  but  the  alcohol,  carbon  dioxide,  and  some  water  escape 
and  puff  up  the  mass  still  more.  The  heat,  however,  soon  hardens  the 
starch,  gluten,  etc.,  into  a  firm  but  porous  loaf. 

Cellulose  (C6H10O5)n  is  widely  distributed  in  the  vegetable 
kingdom.  The  framework  of  all  vegetables  is  cellulose.  It 
is  thus  analogous  to  the  bones  of  animals.  Wood,  cotton, 
linen,  and  paper  are  largely  cellulose.  Pure  cellulose  is  a 
white  substance,  insoluble  in  most  liquids,  but  soluble  in  a 
mixture  of  ammonia  and  copper  oxide.  Concentrated  sul- 
phuric acid  dissolves  it  slowly ;  and  if  the  solution  is  di- 
luted and  boiled,  the  cellulose  is  changed  into  a  mixture 
of  glucose  and  dextrin.  By  this  operation,  wood  could  be 
made  into  a  sugar  and  then  into  alcohol ;  but  the  method 
would  be  too  expensive  to  use  on  a  large  scale. 

Sulphuric  acid  of  a  certain  strength,  if  quickly  and  properly  applied  to 
paper,  changes  it  into  a  tougher  form  called  parchment  paper.  The 
latter  is  often  substituted  for  animal  parchment  (e-g.  sheepskin),  and 
has  a  variety  of  uses. 

Cellulose  has  properties  resembling  those  of  alcohol.  Thus  it  inter- 
acts with  acids  and  forms  ethereal  salts.  With  nitric  acid  it  forms  cellu- 
lose nitrates,  just  as  glycerine  forms  glycerine  nitrates  (see  Nitroglyce- 
rine). The  cellulose  nitrates  are  the  basis  of  smokeless  gunpowders. 
One  of  the  cellulose  nitrates  is  gun  cotton.  It  looks  like  ordinary  cotton, 
and  may  be  spun,  woven,  and  pressed  into  cakes.  It  burns  with  a  large 
flame  if  unconfined ;  but  when  ignited  by  a  percussion  cap  or  when 
burned  in  a  confined  space,  gun  cotton  explodes  violently-  It  is  used  in 
blasting.  Other  cellulose  nitrates  are  known.  Their  solution  in  a  mix- 
ture of  alcohol  and  ether  is  called  collodion.  When  poured  or  brushed 
upon  a  glass  plate  or  the  skin,  the  solvent  evaporates,  leaving  behind  a 
thin  film.  It  is  used  in  preparing  certain  photographic  material  and  as 
a  coating  for  wounds.  The  "  new  skin  "  liquid  recently  offered  for  sale 
is  mainly  collodion.  It  protects  wounds  from  dusty,  impure  air,  and 
thereby  facilitates  the  healing.  A  mixture  of  camphor  and  cellulose  ni- 


Some  Common  Organic  Compounds.        429 

trates  is  called  celluloid.  It  is  easily  molded  into  various  shapes.  The 
white  celluloid  is  made  into  collar  buttons,  and  the  colored  varieties  are 
made  into  toilet  articles  and  ornaments.  Celluloid  smells  of  camphor, 
can  be  lighted  with  a  match,  and  burns  freely  with  a  smoky  flame. 

Paper  is  chiefly  cellulose.  Formerly  it  was  made  from 
various  kinds  of  rags ;  but  now  it  is  made  almost  entirely 
from  wood,  especially  the  paper  used  for  newspapers  and 
cheap  books.  The  best  paper,  such  as  writing  paper,  is 
still  made  from  linen  rags. 

In  making  paper  from  wood,  the  latter  is  reduced  to  a  pulp,  which 
is  washed,  spread  on  a  frame  or  an  endless  wire  gauze,  dried,  and 
pressed.  The  pulp  is  prepared  by  two  processes,  the  mechanical  and 
the  chemical.  Mechanical  pulp  is  made  by  holding  a  stick  of  wood 
against  revolving  stone  upon  which  water  constantly  falls.  Chemical 
pulp  is  made  by  heating  chipped  wood  with  caustic  soda,  or  with  cal- 
cium acid  sulphite  (usually  called  bisulphite).  The  operation  is  con- 
ducted under  pressure  in  huge  tanks  called  digesters.  Chemical  pulp 
has  longer  and  stronger  fibers  than  mechanical  pulp.  The  two  kinds  of 
pulp  are  often  mixed.  Most  paper  is  loaded,  —  that  is,  clay,  gypsum,  or 
other  mineral  matter  is  mixed  with  the  pulp  to  give  the  paper  body. 
Paper  intended  for  printing  or  writing  is  sized,  —  that  is,  the  surface  is 
coated  with  gelatine,  rosin,  or  a  similar  substance  to  prevent  the  ink 
from  spreading.  Many  kinds  are  also  smoothed  by  passing  them 
between  heavy  rollers.  Blotting  and  tissue  papers  are  not  sized  or 
loaded. 

BENZENE    AND    ITS    DERIVATIVES. 

Introduction.  — The  hydrocarbon  benzene  was  mentioned 
in  Chapter  XV  as  the  first  member  of  an  homologous  series. 
In  the  same  chapter  coal  tar  was  described  as  a  black, 
complex  liquid  obtained  as  a  by-product  in  the  manufacture 
of  illuminating  gas.  Now,  coal  tar  is  the  chief  source  of 
benzene  and  some  of  its  related  compounds,  while  from 
benzene  itself  hundreds  of  derivatives  have  been  prepared. 
Some  are  absolutely  indispensable  to  man,  but  many  have 


430  Descriptive  Chemistry. 

as  yet  merely  scientific  interest.     Only  the  most  important 
benzene  compounds  can  be  described  in  this  book. 

Benzene,  C6H6,  is  a  colorless  liquid,  lighter  than  water, 
and  has  an  odor  suggesting  coal  gas.  It  burns  with  a 
luminous,  smoky  flame,  owing  to  its  richness  in  carbon. 
Ordinary  illuminating  gas  owes  its  luminosity  partly  to 
benzene.  It  dissolves  fats,  resins,  iodine,  sulphur,  and 
rubber.  Benzene  is  sometimes  called  benzol.  It  should 
not  be  confused  with  benzine,  which  is  a  mixture  of  hydro- 
carbons derived  from  petroleum.  Benzene  is  chiefly  used 
in  preparing  its  derivatives. 

The  Constitution  of  Benzene  has  been  carefully  studied.  For  rea- 
sons too  extended  to  state  here,  it  is  believed  that  in  a  molecule  of 
benzene  the  carbon  atoms  are  arranged  in  a  ring.  The  structural  for- 
mula is  often  written  thus  — 

H 
I 
C 

/  \ 
H-C  C-H 

II  I 

H-C  C-H 

\    ^ 
C 

I 
H 

Benzene  forms  many  derivatives.  In  all  of  them  the  six  carbon 
atoms  remain  as  a  nucleus.  No  carbon  atom  can  be  removed  from  the 
benzene  molecule  without  producing  complete  decomposition.  But  for 
the  six  hydrogen  atoms,  other  atoms  or  radicals  can  be  substituted. 
Hence,  the  almost  infinite  number  of  derivatives  of  benzene. 

Toluene,  C0H~ .  CH3,  is  the  second  member  of  the  benzene  series. 
It  may  be  regarded  as  methyl  benzene ;  or  as  phenyl  methane,  that  is, 
methane  (CH4)  in  which  one  hydrogen  atom  is  replaced  by  the  radical 
phenyl  (C6H5).  Toluene  is  obtained  from  coal  tar,  and  resembles 
benzene  in  its  properties. 

Nitrobenzene,  C,.H5 .  NO2,  is  a  yellow  liquid  formed  by  the  inter- 
action of  benzene  and  nitric  acid.  It  is  volatile,  and  has  the  odor  of 


Some  Common  Organic  Compounds.        431 

bitter  almonds.  Although  poisonous,  it  is  used  to  produce  the  flavor 
of  almonds  in  essences  and  perfumery.  It  is  chiefly  used,  however,  in 
the  manufacture  of  aniline. 

Aniline,  C6H5.NH2,  is  an  oily  liquid,  slightly  heavier 
than  water.  It  is  prepared  on  a  large  scale  by  reducing 
nitrobenzene  with  nascent  hydrogen.  From  aniline  are 
made  many  compounds  known  as  aniline  dyes.  The 
starting  point  of  these  dyes  is  rosaniline,  which  is  pre- 
pared by  oxidizing  a  mixture  of  aniline  and  toluidine 
(C6H4  .  CH3  .  NH2).  Derivatives  of  rosaniline  produce 
exceedingly  brilliant  colors  in  every  variety  of  shade. 
Vast  dyeing  industries  have  risen  since  the  value  of  coal 
tar  was  discovered  (about  1860). 

Phenol,  C6H5.OH,  is  a  white  crystalline  solid.  It  has 
a  smoky  odor,  is  poisonous,  and  burns  the  skin.  Coal  tar 
is  the  source  of  phenol.  A  solution  of  phenol  in  water, 
popularly  called  carbolic  acid,  is  used  as  a  disinfectant. 

Derivatives  of  Phenol  are  important.  Picric  acid,  or  trinitrophenol 
(C(iH2(NO2)oOH),  is  a  yellow  crystalline  solid  used  in  dyeing  silk  yellow. 
Salts  of  picric  acid  —  the  picrates  —  are  used  in  making  explosives. 
Related  to  phenol  are  hydroquinone  (C6H4(OH)2)  and  pyrogallic 
acid  (C(;Ho(OH)3),  which  are  used  extensively  as  developers  in 
photography. 

Acids,  Aldehydes,  and  Ethereal  Salts  of  the  Benzene 
Series.  —  The  simplest  acid  is  benzoic  acid  (C6H5 .  COOH). 
It  occurs  in  certain  balsams  and  gums.  It  is  usually  pre- 
pared from  gum  benzoin,  and  is  a  white  crystalline  solid 
with  a  fragrant  odor.  The  corresponding  aldehyde  (ben- 
zoic aldehyde,  C6H5.COH)  is  commonly  called  oil  of 
bitter  almonds.  It  is  a  fragrant  liquid  and  is  used  to 
some  extent  as  a  flavoring  substance.  Salicylic  acid 
(C6H4.  OH.  COOH)  is  a  white  crystalline  solid,  which  is 
extensively  used  as  a  food  preservative.  Sodium  salicylate 


432  Descriptive  Chemistry. 

is  a  common  remedy  for  rheumatism.  The  corresponding 
aldehyde  gives  the  fragrance  to  the  wild  flower  known 
as  meadowsweet;  and  methyl  salycilate  is  the  essential 
ingredient  of  the  checkerberry. 

Naphthalene,  C10H8,  is  a  white,  lustrous,  crystalline 
solid  obtained  from  coal  tar.  It  has  a  penetrating,  un- 
pleasant odor,  and  is  used  as  a  substitute  for  camphor 
under  the  name  of  "  moth  balls."  Large  quantities  of 
naphthalene  are  used  in  making  dyestuffs. 

Anthracene,  C14H10,  is  a  white  crystallized  solid,  and, 
like  naphthalene,  is  obtained  from  coal  tar.  It  is  one  of 
the  most  important  hydrocarbons,  because  from  it  alizarin 
is  made.  Alizarin  is  a  valuable  dyestuff,  not  only  because 
it  produces  brilliant  colors  with  different  mordants,  but 
also  because  most  of  these  colors  are  fast,  that  is,  they 
do  not  fade  like  many  aniline  colors.  The  Turkey  red -so 
common  on  cotton  goods,  is  produced  by  alizarin.  Aliza- 
rin was  formerly  obtained  from  madder  root,  but  now  vast 
quantities  are  artificially  prepared. 

Glucosides  are  substances  occurring  in  many  plants  and  vegetables. 
By  the  action  of  ferments  they  are  changed  into  glucose  and  other 
substances  that  are  benzene  derivatives.  Amygdalin,  for  example,  is 
found  in  bitter  almonds,  cherry  and  peach  kernels,  and  laurel  leaves. 
The  ferment  emulsin,  which  also  occurs  in  the  plants,  breaks  up  the 
amygdalin  into  oil  of  bitter  almonds,  hydrocyanic  acid,  and  glucose. 
Tannin  is  also  a  glucoside.  The  tannins  are  a  group  of  related  com- 
pounds found  in  the  leaves,  bark,  and  other  parts  of  the  oak,  hemlock, 
and  pine  trees,  in  sumach,  gallnuts,  tea,  coffee,  and  numerous  plants. 
Several  acids  have  been  obtained  from  tannins.  The  best  known  are 
gallic  acid  and  tannic  acid ;  the  latter  is  also  often  called  simply  tan- 
nin, and  probably  all  tannins  contain  some  tannic  acid.  Tannic  acid 
changes  into  gallic  acid  according  to  the  following  equation  — 

C14H1009         +         H20  2C7Hfi05 

Tannic  Acid  Gallic  Acid 


Some  Common  Organic  Compounds.        433 

The  formula  of  gallic  acid  may  be  written  C(;H2  (OH)3  .  COOH,  thus 
showing  its  relation  to  benzene.  Tannin,  in  whatever  form,  produces 
black  compounds  with  iron  salts.  Its  presence  in  tea,  hemlock  bark, 
etc.,  may  be  shown  by  the  formation  of  a  black  precipitate  upon  the 
addition  of  ferrous  sulphate.  This  property  is  utilized  in  making 
writing  ink,  though  some  kinds  of  ink  are  now  made  from  aniline 
dyes.  The  tannin  in  oak  and  hemlock  barks  is  used  in  tanning  leather. 
When  raw  hides  are  soaked  in  solutions  of  tannin,  the  tannic  acid 
changes  certain  substances  in  the  skin  into  insoluble  compounds, 
which  remain  in  the  hide,  thereby  converting  it  into  the  soft  pliable 
form  known  as  leather.  Tannins  are  also  used  as  mordants  in  dyeing 
silk,  cotton,  and  linen. 

Alkaloids  are  complex  compounds  obtained  from  plants  and  vegeta- 
bles. The  chief  property  is  the  power  to  produce  marked  physiological 
effects  upon  animals.  All  of  them  contain  nitrogen,  and  resemble 
ammonia  in  having  an  alkaline  reaction  and  in  uniting  directly  with 
acids  to  form  salts.  Their  commercial  form  is  usually  a  salt.  Many 
are  used  as  medicines  and  drugs,  although  they  are  poisonous,  especially 
if  taken  in  large  quantities.  Theine  or  caffeine  is  the  alkaloid  obtained 
from  tea  and  coffee.  Nicotine  comes  from  tobacco  and  is  very  poison- 
ous. Cocaine  is  obtained  from  the  coca  plant.  One  of  its  salts  is 
used  by  surgeons  and  dentists  to  relieve  pain.  Quinine  and  cinchonine 
are  extracted  from  the  bark  of  the  cinchona  tree ;  both  are  used  as  a 
remedy  for  fevers.  Morphine  is  the  chief  alkaloid  found  in  opium.  The 
latter  is  the  dried  sap  obtained  from  a  certain  part  of  the  unripe  poppy. 
Morphine  in  different  forms  is  used  to  relieve  pain  and  induce  sleep. 
The  two  familiar  medicines,  laudanum  and  paregoric,  contain  prepara- 
tions of  opium.  Large  doses  of  any  form  of  opium  may  be  fatal. 

EXERCISES. 

1.  How  were  organic  and  inorganic  compounds'  once  defined  ?    Do 
they  differ  fundamentally  ?     What  compounds  are  now  included  by  the 
term  organic? 

2.  What  is  the  essential  element  in  organic  compounds  ?     What 
other  elements  are  often  present  ? 

3.  Give  four  reasons  for  the  vast  number  of  organic  compounds. 

4.  Define  an  organic  radical.     Name  three. 

5.  Define  constitution.     Illustrate  it  by  the  empirical,  rational,  and 
graphic  formulas  of  alcohol. 


434  Descriptive  Chemistry. 

6.  Name  the  nine  important  groups  of  organic  compounds. 

7.  Review  the  general  properties  of  hydrocarbons  (see  Chapter  XV) . 
Name  four  hydrocarbons. 

8.  Define  an  alcohol.     Discuss  the  constitution  of  alcohols. 

9.  Describe  the  preparation  of  methyl  alcohol.     State  its  properties 
and  uses.     Why  is  it  called  (a}  methyl  alcohol,  and  (b)  wood  alcohol  ? 

10.    State  (a)  the  properties,  and  (b)  the  uses  of  ethyl  alcohol. 
n.    What  is  (a)  alcohol,   (b)  ethyl   alcohol,   (c)  absolute   alcohol, 
(d)  methylated  spirit,  (e)  proof  spirit  ? 

12.  What  is  fermentation  ?     What  are  ferments  ? 

13.  Describe  the  preparation  of  alcohol.     Discuss  the  preparation, 
composition,  and  properties  of  («)  wines  and  beers,  and  (<£)  distilled 
liquors. 

14.  What  are  aldehydes  ?     How  are  they  related  to  alcohols  and  to 
hydrocarbons  ? 

15.  Describe  the  preparation  and  properties  of  (a}  acetic  aldehyde, 
and  (b)  formic  aldehyde.     State  the  uses  of  the  latter.     What  is  its 
commercial  name  ? 

1 6.  What  are  ethers  ?     How  are  they  related  to  alcohols  ? 

17.  Describe  the  preparation,  and  state  the  properties  and  uses  of 
ordinary  ether. 

1 8.  What  are  organic  acids  ?     Illustrate  (by  acetic  acid)  their  rela- 
tion to  hydrocarbons,  alcohols,  and  aldehydes. 

19.  Describe  the  manufacture  of  acetic  acid.    State  (#)  its  properties, 
and  (£)  its  uses. 

20.  What  is  (a)  pyroligneous  acid,  (£)  glacial  acetic  acid,  (<:)  wood 
vinegar,  (d)  commercial  acetic  acid  ? 

21.  Discuss  the  composition  of  acetic  acid. 

22.  What  is  vinegar  ?     Describe  its  manufacture.     State  its  proper- 
ties and  uses. 

23.  What  are  acetates  ?     State  their  general   properties.     Describe 
four,  and  state  their  uses. 

24.  Name  three  other  acids  (besides  acetic)  in  the  fatty  acid  series. 
Why  is  this  series  so  called  ? 

25.  State  the  occurrence,  properties,  and  uses  of  (a*)  oxalic  acid, 
(£)  lactic  acid,  (c)  tartaric  acid,  (d}  citric  acid.     Where  is  malic  acid 
found  ? 

26.  What  is  (a}  argol,  (b)  crude  tartar,  (c)  cream  of  tartar,  (d)  tar- 
tar emetic  ? 


Some  Common  Organic  Compounds.        435 

27.  Review  baking  powder  (see  Sodium  Bicarbonate). 

28.  What  are  ethereal  salts  ?     How  are  they  formed  ?     Where  are 
they  found  ?     Describe  ethyl  acetate.     Name  three  other  ethereal  salts 
and  state  their  properties. 

29.  What  is  the  test  for  (a)  alcohol,  and  (£)  acetic  acid  ? 

30.  State  clearly  the  general  relations  of  fats  to  glycerine  and  soap. 

31 .  Name  the  chief  ingredients  of  fats  and  oils.     What  is  (a)  tallow, 
(b}  butter,  (c}  oleomargarine,  (d)  stearin  ? 

32.  Describe  the  preparation  of  glycerine.     State  its  properties  and 
uses. 

33.  Discuss  the  constitution  of  glycerine.     State  the  properties  and 
uses  of  (a)  nitroglycerine,  and  (£)  dynamite. 

34.  What  is  soap  ?     Describe  its  general  method  of  manufacture. 
What  is  the  chemistry  of  its  manufacture  ?     What  fats  and  alkalies  are 
used  in  making  soap  ?    Describe  (a)  the  cold  process,  and  (b)  the  boil- 
ing process  of  soap  making. 

35.  What   are   carbohydrates?     Why  is   this   term   used?     Name 
several  carbohydrates. 

36.  What  are  sugars  ?     Name  several. 

37.  Discuss  the  distribution  of  cane  sugar.     State  its  properties. 
What  is  (a)  cane  sugar,  (b)  sucrose,  (c}  saccharose,  (d)  barley  sugar, 
(e)  caramel  ?     For  what  is  the  last  used  ? 

38.  Describe  the  preparation  of  raw  sugar  from  (a)  sugar  cane,  and 
(b)  sugar  beets. 

39.  Describe  the  refining  of  sugar. 

40.  What  is  (a)  granulated  sugar,  {b}  brown  sugar,  (c}  molasses  ? 

41.  What  is  the  sugar  of  milk  ?     What  is  its  scientific  name  ?     For 
what  is  it  used  ? 

42.  What  is  the  formula  of  glucose  ?     WThat  other  names  has  glu- 
cose ?     Where  is  glucose  found  ?     What  sugar  is  closely  related  to 
glucose  ?    How  is  glucose  formed  from  cane  sugar  ?    State  the  equation 
for  the  reaction. 

43.  How  is  commercial  glucose  prepared  ?     What  is  (a)  commercial 
grape  sugar,  and  (b}  "  glucose "  ?     State  the  properties  and  uses  of 
commercial  glucose. 

44.  Describe  the  test  for  sugar. 

45.  Discuss  the  distribution  of  starch.     Describe  starch.     State  its 
properties.     What  is  the  test  for  starch  ? 

46.  How  is  starch  prepared  ?     State  its  uses. 


436  Descriptive  Chemistry. 

47.  What  is  the  simplest  formula  of  starch  ?     How  does  it  differ 
from  the  formula  of  («)  cane  sugar,  and  ($)  glucose  ? 

48.  What  is  dextrin  ?     How  is  it  prepared  ?     For  what  is  it  used  ? 

49.  Discuss  the  chemistry  of  bread  making. 

50.  What  is  cellulose  ?    Describe  pure  cellulose.    State  its  properties. 

51.  What  is  (a)  parchment  paper,  (£)  gun  cotton,  (c)  collodion  ? 

52.  What  is  the  chief  constituent  of  paper  ?     Describe  the  manufac- 
ture of  paper. 

53.  State  the  source  of  benzene.     State  its  properties.     What  is  (a) 
benzol,  and  (b)  benzine  ? 

54.  To  what   class   of  organic  compounds  does   benzene   belong? 
Why  is  it  such  an  important  compound  ? 

55.  What  is  the  chemical  relation  of  benzene  to  (a)  toluene,  (£) 
nitrobenzene,  (c)  aniline,  (d}  phenol,  (e)  benzole  acid  ? 

56.  Describe  nitrobenzene.     What  is  its  chief  use  ? 

57.  Describe  aniline.     How  is  it  prepared  ?     For  what  is  it  used  ? 

58.  Describe  phenol.     What  is   its  source  and  use  ?     What  is  its 
common  name  ? 

59.  State  briefly  the  relation  of  phenol  to  (a)  picric  acid,  (]£)  pi- 
crates,  (c}  hydroquinone,  (//)  pyrogallic  acid.     What  is  the  use  of  each  ? 

60.  Describe  briefly  benzoic  acid  and  benzoic  aldehyde. 

61.  Describe  salicylic  acid.     State  the  use  of  this  acid. 

62.  Describe  naphthalene.    What  is  its  popular  name  ?    State  its  uses. 

63.  Describe  anthracene.     State  its  use.     What  is  alizarin  ? 

64.  What  are  glucosides  ?     Discuss  (a}  the  occurrence,  (£)  the  prop- 
erties, and  (c)  the  uses  of  tannin.     What  is  (a)  ink,  and  (£)  leather  ? 

65.  What  are  alkaloids  ?     Name  six.     What  is  their  chief  property  ? 

PROBLEMS. 

1.  Alcohol  is  0.8  as  heavy  as  water.     What  is  the  weight  of  1200  cc. 
of  alcohol  ? 

2.  If  10  gm.  of  pure  alcohol  are  burned,  what  weight  of  each  product 
is  formed  ?     (Equation  is  C2H(;O  +  30,  =  2  CO2  +  3  H2O.) 

3.  Calculate  the  percentage  composition  of  (#)  alcohol  (C2H6O),  (£) 
acetic  acid  (C2H4O2,  (c}  cane  sugar  (C]2H2,On). 

4.  Calculate  the  simplest  formulas  of  the  substances  having  the  com- 
position :     (a}     carbon  =  40,    hydrogen  =  6.67,     oxygen  =  53.33  ;     (£) 
carbon  =  15.8,  hydrogen  =  5.26,  nitrogen  =  36.84,  sulphur  =  42.1  ;   (c) 
carbon  =  54.55,  hydrogen  =  9.09,  oxygen  =  36.36. 


APPENDIX. 


1.  The  Metric  System.  —  The  fundamental  unit  of  this  system  of 
weights  and  measures  is  the  meter.  It  is  the  unit  of  length,  and  is 
39.37  inches  long. 

The  meter  and  the  other  units  have  multiples  and  submultiples, 
which  are  designated  by  prefixes  attached  to  the  particular  unit.  The 
multiple  prefixes  are  deca-,  hecto-,  and  kilo-,  equivalent  respectively  to 
10,  100,  and  1000.  The  submultiple  prefixes  are  deci-,  centi-,  and  milli-, 
which  correspond  respectively  to  o.i,  o.oi,  and  o.ooi. 

The  unit  of  weight  is  the  gram.  It  is  derived  from  the  kilogram, 
which  is  the  weight  of  a  cubic  decimeter  of  water  at  4°  C.  A  kilogram 
weighs  about  2.2  pounds.  Small  weights  are  expressed  in  terms  of  the 
gram.  Thus,  the  weight  of  an  object  weighing  2  grams,  2  centi- 
grams, and  5  milligrams  is  2.025  grams. 

The  unit  of  volume  is  the  liter.  It  is  equal  to  the  capacity  of  the 
vessel  containing  a  kilogram  of  water.  A  liter  equals  about  one  quart. 

The  relation  between  the  units,  multiples,  and  submultiples  is  shown 
in  the  — 

TABLE  OF  THE  METRIC  SYSTEM. 


LENGTH. 

WEIGHT. 

VOLUME. 

NOTATION. 

Kilometer 

Kilogram 

Kiloliter 

1000. 

Hectometer 

Hectogram 

Hectolitefr 

100. 

Decameter 

Decagram 

Decaliter 

10. 

METER 

GRAM 

LITER 

I. 

Decimeter 

Decigram 

Deciliter 

O.I 

Centimeter 

Centigram 

Centiliter 

0.01 

Millimeter 

Milligram 

Milliliter 

O.OOI 

From  this  table  it  is  evident  that  10  milligrams  equal  I  centigram,  10 
centigrams  equal  I  decigram,  10  decigrams  equal  i  gram,  and  so  on. 

4.37 


438 


Descriptive  Chemistry. 


The  relation  of  the  metric  system  to  weights  and  measures  in  com- 
mon use  is  shown  by  the  — 

TABLE  OF  METRIC  EQUIVALENTS. 


meter          =  39.37  inches 
kilometer   =  0.62  mile 
centimeter  =  0.39  inch 
liter  =  0.908  quart 

liter  =  1.056  quart '(liq.) 

gram  =  15.432  grains 

kilogram     =  2.2  pounds  (avoir.) 
metric  ton  =  2204  pounds 


inch 
mile 

cubic  inch 
quart  (liq.) 
pound  (avoir.) 
ounce  (avoir.) 
ounce  (troy) 
grain  (apoth.) 


2.54  centimeters 

1.6  kilometers 

16.39  cubic  centimeters 

0.9465  liter 

0.4536  kilogram 

28.35  grams 

31.1  grams 

0.0648  gram 


The  passage  from  the  English  to  the  metric  system  may  be  accom- 
plished by  utilizing  the  — 

TABLE  OF  METRIC  TRANSFORMATION. 


To  CHANGE 


MULTIPLY  BY 


Inches  to  centimeters    
Centimeters  to  inches    

2-54 
0-3937 
16.387 

Cubic  centimeters  to  cubic  inches        

0.061 
28.31; 

Grams  to  ounces  (avoir.)       

0^0353 
0.0648 

Grams  to  grains      

J543 

The  customary  abbreviations  of  the  common  denominations  are  — 


meter,  m. 
decimeter,  dm. 
centimeter,  cm. 


liter,  1. 

kilogram,  kg.  or  Kg. 

decigram,  dg. 


cubic  centimeter,  cc. 
milligram,  mg. 
centigram,  eg. 


The  ^referable  abbreviation  for  gram  is  gm.     The  same  abbreviation 
is  used  for  singular  and  plural,  e.g.  I  m.,  4  gm.,  3  cm.,  50  cc. 

A  convenient  relation  (true  only  in  the  case  of  water)  to  remember 
is  i  I  =  i  kg.  =  i  cu.  dm.  =  1000  cc-  ^  1000  gm.  =  2,3  lb. 


Appendix. 


439 


PROBLEMS. 

1.  What  is  the  abbreviation  of  gram,  centigram,  liter,  meter,  cubic 
centimeter,  centimeter,  decimeter,  milligram  ? 

2.  Express  (a)  i  liter  in  cubic  centimeters,  (£)  2  1.  in  cc.,  (c)  i  meter 
in  centimeters,  (d)  250  cm.  in  dm.,  (e)   i  kg.  in  grams,  (/)  250  gm. 
in  mg. 

3.  Add  2  kg.,  5  dg.,  2  eg.,  4  gm.,  and  7  mg.,  and  express  the  sum  in 
grams. 

4.  How  many  cc.  in  a  liter  ? 

5.  What  is  the  weight  in  grams  of  (a)  i  liter  of  water,  (<£)  250  cc., 
(c)  500  cc.,  (d)  721  cc.  ? 

6.  Express  in  grams  (a)  721  kg.,  (£)  62  mg.,  (c)  245  eg.,  (d}  84  dg. 

7.  Express   (a)  40  meters  in  inches,  (£)  25  kilograms  in  pounds, 
(c}  54  grams  in  ounces,  (d)  72  grams  in  grains,  (e)  75  liters  in  quarts 
(liq.). 

2.  The  Thermometer  in  scientific  use  is  the  centigrade.  The  boil- 
ing point  of  water  on  this  thermometer  is  zoo,  and  the  freezing  point  is 
o  (Fig.  80).  The  equal  spaces  between  these  points  are  called  degrees. 
The  abbreviation  for  centigrade  is  C.,  and  for  degrees 
is  °.  Thus,  the  boiling  point  of  water  is  100°  C. 
Degrees  below  zero  are  always  designated  as  minus, 
e.g.  —12°  C.,  means  12  degrees  below  zero. 

The  thermometer  in  popular  use  is  the  Fahrenheit. 
On  this  instrument  the  boiling  point  of  water  is  212° 
and  the  freezing  point  is  32°  above  zero  (Fig.  80). 

To  change  Fahrenheit  degrees  into  the  equivalent 
centigrade  degrees,  subtract  32  and  multiply  the 
remainder  by  f ,  or  briefly  — 

C  =  f(F-32). 


To  change  centigrade  degrees  into  the  equivalent 
Fahrenheit  temperature,  multiply  by  f  and  add  32  to 
the  product,  or  briefly — 


212 


FIG.  80.  — Ther- 
mometers. 


The  point  —  273°  C.  is  called  absolute  zero.  Absolute  temperature 
is  reckoned  from  this  point.  Degrees  on  the  absolute  scale  are  found 
by  adding  273  to  the  readings  on  the  centigrade  thermometer.  Thus, 
273°  absolute  is  o°  C.,  274°  absolute  is  + 1°  C.,  etc. 


440 


Descriptive  Chemistry. 


PROBLEMS. 

1.  Change  into  Fahrenheit  readings  the  following  centigrade  read- 
ings :   (a)  60.5,  (J)  40,  (0  92,  (<0  -  5>  (')  o,  (/)ioo,  (£)  860,  (A)  -40. 

2.  Change  into  centigrade  readings  the  following  Fahrenheit  read- 
ings:   (a)  207,  (b)  1 80,  (0  o,  (W)  -30,  (*)  212,  (7)   100,  (£)   -40, 
(X)  270. 

3.  Express  the  following  centigrade  readings  in  absolute  readings : 
(«)  o,  (*)  24,  (*)  -13,00  -26°- 

3.  Crystallization.  —  Most  substances  in  passing  from  a  liquid  or  a 
gas  into  a  solid  assume  a  definite  shape.  This  change  is  called  crys- 
tallization, and  the  substances  are  said  to  crystallize  or  to  form  crys- 
tals. Crystals  are  produced  by  (i)  evaporating  a  solution,  (2)  cooling 
a  melted  solid,  or  (3)  cooling  a  vapor.  Thus,  salt  crystals  are  formed 
by  evaporating  a  salt  solution ;  sulphur  crystals,  by  melting  and  then 
cooling  sulphur,  and  iodine  crystals,  by  heating  iodine  in  a  test  tube. 
These  methods  are  called,  respectively,  evaporation,  fusion,  and  sub- 
limation. 

As  a  rule  each  substance  has  an  individual  crystal  form  by  which  it 
can  be  distinguished.  Although  there  are  thousands  of  different  crys- 
tals, all  belong  to  one  of  six  classes  or  systems.  This  classification  is 
based  upon  two  assumptions:  (i)  all  crystals  contain  certain  lines 
called  axes,  and  (2)  the  surfaces  or  faces  are  grouped  around  the  axes 
in  definite  positions.  The  axes  connect  angles,  edges,  or  faces,  which 
are  similarly  situated  on  opposite  sides  of  the  crystal.  The  bounding 
planes  or  faces  are  arranged  symmetrically  around  the  axes,  which  also 
determine  (by  their  lengths  and  relative  positions)  the  positions  of  the 
bounding  planes.  For  example,  the  cube  has  three  equal  axes  at  right 
angles  to  one  another,  and  terminating  in  the  center  of  each  of  the  six 
bounding  surfaces. 

The  following  is  a  brief  description  of  the  six  'systems  of  crystal- 
lization :  — 


FIG.  81.  —  Isometric  crystals  (cube,  octahedron,  dodecahedron). 


Appendix. 


441 


(i)  Isometric.  —  This  has  three  equal  axes  intersecting  at  right 
angles.  The  simplest  forms  are  the  cube,  octahedron,  and  dodecahedron 
(Fig.  81).  Substances  crystallizing  in  this  system  are  diamond,  com- 
mon salt,  alum,  fluor  spar,  iron  pyrites,  and  garnet. 


FIG.  82.  — Tetragonal  crystals. 

(2)  Tetragonal.  —  This  has  three  axes  at  right  angles  ;  but  one  axis 
is  shorter  or  longer  than  the  other  two,  which  are  equal.  The  common 
forms  are  the  prism,  pyramid,  and  their  combinations  (Fig.  82).  Tin 
dioxide  and  zircon  form  tetragonal  crystals. 


FIG.  83.  —  Orthorhombic  crystals. 

(3)  Orthorhombic.  —  This  has  three  unequal  axes  intersecting  at 
right  angles.  Common  forms  are  the  prism,  pyramid,  and  their  com- 
binations (Fig.  83).  Potassium  nitrate,  barium  sulphate,  topaz,  and 
native  sulphur  crystallize  in  this  system  (see  Fig.  49). 


FIG.  84.  —  Hexagonal  crystals. 


(4)  Hexagonal.  —  This  has  four  axes  :  three  are  equal  and  intersect 
at  60°  in  the  same  plane  ;  the  fourth  is  longer  or  shorter  than  the  others 


442 


Descriptive  Chemistry. 


and  is  at  right  angles  to  their  plane.  It  is  a  complex  system.  Common 
forms  are  the  prism,  pyramid,  rhombohedron,  scalenohedron,  and  their 
combinations  (Fig.  84).  In  this  system  are  found  quartz,  calcite,  beryl, 
corundum,  and  ice  (see  Figs.  5,  52,  61). 

(5)  Monoclinic.  —  This  has  three  unequal  axes  :  two  cut  each  other 
obliquely,  and  the  third  is  at  right  angles  to  the  plane  of  the  other  two. 
Common  forms  are  combinations  of  prisms.  It  is  a  complex  system, 
but  includes  many  substances,  e.g.  sulphur  deposited  by  fusion,  sodium 
carbonate,  borax,  gypsum,  and  ferrous  sulphate  (Fig.  85). 


FIG.  85.  —  Monoclinic  crystal. 


FIG.  86.  —  Triclinic  crystals. 


(6)  Triclinic.  —  This  has  three  unequal  axes,  all  intersecting  at 
oblique  angles.  Common  forms  are  complex  combinations.  Copper 
sulphate,  potassium  dichromate,  boric  acid,  and  several  minerals  form 
triclinic  crystals  (Fig.  86). 

4.  History  and  Biography.  —  The  biographical  data  and  table 
given  here  will  serve  as  a  basis  for  this  interesting  branch  of  chemistry. 
Additional  facts  can  be  obtained  from  the  historical  books  mentioned 
below  (under  "  Reference  Books  ") . 

Arrhenius,  Svante,  1859  .  Swedish  physicist.  Contributor  to 

modern  theory  of  solution. 

Avogadro,  Amadeo,  1776-1856.  Italian  chemist  and  physicist.  Pro- 
posed in  1811  his  hypothesis  —  equal  number  of  molecules  in  equal 
volumes  of  all  gases  at  same  temperature  and  pressure. 

Balard,  Antoine  Jerome,  1802-1876.  French  chemist.  Discovered 
bromine  in  1826. 

Becher,  Johann  Joachim,  1635-1682.  German  physician.  Dis- 
covered few  facts,  but  collected  and  explained  writings  of  others. 
Believed  in  alchemy,  but  made  no  search  for  gold.  Laid  foundations  of 
phlogiston  theory. 


Appendix.  443 

Bergman,  Torbern,  1735-1784.  Swedish  chemist.  Improved 
methods  of  chemical  analysis.  Believed  in  phlogiston.  Studied  min- 
erals and  organic  acids.  Contributed  much  to  the  industrial  develop- 
ment of  Sweden.  Intimate  friend  of  Scheele. 

Berthollet,  Claude  Louis,  1748-1822.  French  chemist.  Studied 
composition  of  ammonia,  properties  and  nature  of  chlorine,  hydrogen 
sulphide,  and  hydrocyanic  acid.  Explained  chemical  changes  by 
"  affinity."  His  discussion  with  Proust  led  to  law  of  definite  proportions. 

Berzelius,  Johann  Jacob,  1779-1848.  Swedish  chemist.  Deter- 
mined many  atomic  weights.  Introduced  use  of  symbols.  Discovered 
selenium,  prepared  silicon  and  several  rare  elements.  Investigated  law 
of  multiple  proportions,  proposed  dualistic  theory  and  an  electrochem- 
ical theory,  improved  experimental  methods.  Industrious  investigator, 
prolific  writer. 

Bessemer,  Sir  Henry,  1813-1898.  English  metallurgist.  Devised, 
in  1856,  Bessemer  process  of  making  steel. 

Black,  Joseph,  1728-1799.  Scotch  chemist  and  physicist.  Dis- 
covered carbon  dioxide.  Showed  relation  of  this  gas  to  carbonates  of 
alkalies  and  alkaline  earths.  Opposed  phlogiston  theory.  Teacher 
and  friend  of  James  Watt  and  Rutherford. 

Boyle,  Robert,  1626-1691.  English  philosopher.  Announced  law 
of  effect  of  pressure  on  gases.  Studied  air  and  water.  Opposed 
to  alchemy.  Views  anticipated  present  conception  of  constitution  of 
matter.  Laid  foundation  of  qualitative  analysis. 

Bunsen  von,  R.  W.  E.,  1811-1899.  German  chemist.  Studied 
blast  furnace  and  developed  gas  analysis.  Invented  the  burner,  pho- 
tometer, and  battery  bearing  his  name.  With  Kirchhoff  (about  1860) 
devised  the  spectroscope,  and  by  it  developed  spectrum  analysis  and 
discovered  rubidium  and  caesium ;  improved  the  calorimeter ;  studied 
chemical  action  of  light. 

Cannizzaro,  Stanislao,  1826 .  Italian  chemist.  Revived  Avoga- 

clro's  hypothesis  in  1858,  and  thereby  led  to  revision  of  atomic  weights. 

Cavendish,  Henry,  1731-1810.  English  chemist.  Discovered  hy- 
drogen, determined  specific  gravity  of  gases,  showed  (i)  solubility  of 
calcium  carbonate  in  water  containing  carbon  dioxide,  (2)  formation  of 
water  by  burning  of  hydrogen.  Determined  composition  of  the  atmos- 
phere and  of  nitric  oxide.  Accepted  phlogiston  theory.  He  was 
parsimonious,  eccentric,  shy ;  trained  mathematician  and  electrician ; 
"  the  richest  of  the  wise,  and  the  wisest  of  the  rich." 


444  Descriptive  Chemistry. 

Charles,  Jacques  Alex  Cesar,  1746-1822.  French  physicist.  Pro- 
posed law  bearing  his  name. 

Courtois,  Bernard,  1777-1838.  French  chemist.  Discovered  iodine 
in  1811. 

Dalton,  John,  1766-1844.  English  chemist,  physicist,  and  mathe- 
matician. Devised  atomic  theory.  Discovered  law  of  multiple  propor- 
tions. "  Dalton  was  often  inaccurate  as  to  facts,  deficient  in  the  details 
of  chemical  manipulations,  and  did  not  hold  high  rank  as  an  experi- 
menter; but  he  was  good  at  drawing  conclusions  and  at  stating 
generalizations,  his  aim  being  the  establishment  of  general,  underlying 
laws."  (Venable.) 

Davy,  Sir  Humphry,  1778-1829.  English  chemist.  Studied  gases, 
demonstrated  properties  of  nitrous  oxide,  determined  composition  of 
hydrochloric  acid,  studied  iodine  and  chlorine,  named  latter.  Isolated 
potassium,  sodium,  barium,  calcium,  and  strontium  by  electrolysis,  and 
studied  action  of  electricity  on  water  and  on  many  other  substances. 
Devised  miner's  safety  lamp.  "He  was  one  of  the  most  brilliant 
chemists  the  world  has  ever  seen  and  the  greatest  England  has  pro- 
duced." 

Dewar,  James,  1842 .  English  chemist.  Pioneer  in  the  lique- 
faction of  gases  by  modern  methods.  (See  Hydrogen.) 

Dulong,  Pierre  Louis,  1785-1838.  French  chemist  and  physicist. 
With  Petit  announced  law  of  specific  heats  in  1819. 

Dumas,  Jean  Baptiste  Andre,  1800-1884.  French  chemist.  Deter- 
mined many  atomic  weights,  gravimetric  composition  of  water,  compo- 
sition of  air.  Investigated  many  organic  compounds.  Devised  a 
method  of  determining  vapor  density.  Excellent  teacher,  careful 
editor,  and  faithful  public  servant. 

Faraday,  Michael,  1791-1869.  English  chemist  and  physicist. 
Liquefied  chlorine  and  other  gases.  Showed  quantitative  relation  be- 
tween electric  current  and  chemical  changes,  and  developed  electro  chem- 
istry. Was  Davy's  assistant  and  successor  in  the  Royal  Institution. 
Popular  lecturer,  keen  investigator,  and  ardent  lover  of  science. 

Gay-Lussac,  Joseph  Louis,  1778-1850.  French  chemist  and  physi- 
cist. Announced  law  of  gas  volumes  in  1808.  Worked  on  cyanogen, 
iodine,  halogen  acids,  alkaline  oxides,  isolation  of  boron.  Improved 
methods  of  analyzing  organic  compounds.  Was  pupil  of  Berthollet. 
"  Was  a  trained  chemist,  capable  of  most  accurate  analytical  work, 
and  possessing  scientific  acumen  in  a  very  high  degree."  (Venable.) 


Appendix.  445 

Glauber,  Johann  Rudolph,  1604-1668.  German  chemist.  Believed 
in  alchemy.  Discovered  sodium  sulphate,  which  even  now  bears  his 
name.  Suggested  improvements  in  industrial  chemistry. 

Graham,  Thomas,  1805-1869.  British  chemist.  Studied  diffusion 
of  gases,  acids  of  phosphorus,  water  of  crystallization,  and  dialysis. 
Developed  idea  of  basicity  of  acids. 

Hofmann  von,  August  Wilhelm,  1818-1892.  German  chemist. 
Studied  organic  chemistry  exhaustively.  Coal-tar  industry  arose 
largely  from  his  work.  Devised  unique  lecture  apparatus,  e.g.  that  for 
the  electrolysis  of  water.  Brilliant  teacher,  prolific  investigator. 

Kirchhoff ,  Gustav  Robert,  1 824-1 887.  German  physicist.  With  Bun- 
sen,  devised  spectroscope  and  founded  principles  of  spectrum  analysis. 

Lavoisier,  Antoine  Laurent,  1743-1794.  French  chemist.  Over- 
threw phlogiston  theory,  explained  combustion,  contributed  many  facts 
to  a  large  number  of  chemical  topics.  Devised  foundation  of  chemical 
nomenclature.  Interpreted  experiments  of  other  chemists.  Efficient 
public  servant.  Regarded  by  many  as  the  founder  of  modern  chem- 
istry. Accused  of  appropriating  public  money  and  of  "  putting  water 
in  the  people's  tobacco,"  he  was  condemned  by  the  infamous  Robes- 
pierre, and  publicly  guillotined. 

Liebig  von,  Justus,  1803-1873.  German  chemist.  Laid  founda- 
tions of  agricultural  and  organic  chemistry.  Eminent  teacher. 

Mendeleeff,  Dmitri  Ivanovitch,  1834 .  Russian  chemist.  An- 
nounced periodic  law  in  1868. 

Meyer,  Lothar  1830-1895.  German  chemist.  Contributed  to  estab- 
lishment of  periodic  law. 

Moissan,  Henri,  1852 .  French  chemist.  Isolated  fluorine, 

devised  and  perfected  electric  furnace,  prepared  artificial  diamonds, 
rare  metals,  and  refractory  compounds. 

Ostwald,  Wilhelm,  1853 .  German  chemist.  Contributor  to 

modern  theory  of  solution.  Eminent  teacher  and  prolific  writer. 

Petit,  Alexis  Therese,  1791-1820.     French  physicist.    (See  Dulong.) 

Priestley,  Joseph,  1733-1804.  English  chemist  and  theologian. 
Student  of  electricity,  light,  and  gases.  Discovered  oxygen.  Devised 
pneumatic  trough.  His  political  and  religious  views  were  so  freely 
expressed  that  he  was  obliged  to  leave  England.  Came  to  America  in 
1795.  Died  at  Northumberland  near  Philadelphia,  Pennsylvania. 

Proust,  Louis  Joseph,  1755-1826.  French  chemist.  Defended 
law  of  definite  proportions  in  a  long  controversy  with  Berthollet. 


446  Descriptive  Chemistry. 

"  One  of  the  good  results  of  this  controversy  was  to  bring  about  a  defi- 
nition of  compounds  and  mixtures,  and  a  clear  distinction  between 
them.  In  course  of  it,  also,  Proust  discovered  the  hydroxides,  a  class 
of  compounds  until  then  confused  with  the  oxides."  (Venable.) 

Prout,  William,  1785-1850.  English  physician.  Advanced  in  1815 
the  hypothesis  that  the  atomic  weights  of  all  elements  are  whole 
numbers. 

Ramsay,  William,  1852  .  English  chemist.  Discovered 

argon,  helium,  neon,  krypton,  and  xenon. 

Rutherford,  Daniel,  1749-1819.  Scotch  botanist  and  physician. 
Discovered  nitrogen  in  1772.  Pupil  of  Black. 

Scheele,  Carl  Wilhelm,  1742-1786.  Swedish  chemist.  Discovered 
chlorine,  ammonia,  manganese,  baryta,  many  acids  (organic  and  inor- 
ganic), and  oxygen  (independently  of  Priestley).  Isolated  and  studied 
borax,  glycerine,  Prussian  blue,  microcosmic  salt.  Improved  the 
methods  of  preparing  many  substances.  Was  very  poor.  Friend  and 
companion  of  Bergman.  Achieved  marvelous  results  with  simple 
appliances.  Believed  in  phlogiston. 

Stahl,  George  Ernst,  1660-1734.  German  physician  and  chemist. 
Revived  and  extended  Becher's  ideas  of  combustion.  Introduced  the 
name  phlogiston.  Strongly  advocated  this  theory.  Successful  teacher 
and  writer. 

Stas,  Jean  Servais,  1813-1891.  Belgian  chemist.  Determined 
accurately  many  atomic  weights.  Pupil  of  Dumas.  Overthrew  Prout's 
hypothesis. 

Van  Helmont,  Jean,  1577-1644.  Dutch  chemist.  Studied  gases, 
and  discovered  carbon  dioxide.  Had  imperfect  but  introductory  views 
on  physiological  chemistry,  indestructibility  of  matter,  and  elements. 
Believed  in  the  alkahest  or  universal  solvent. 

Van't  Hoff,  Jacobus  Hendricus,  1852 .  Dutch  chemist.  Con- 
tributor to  chemistry  of  space  relations  of  atoms  and  to  modern  theory 
of  solution. 

Wohler,  Friedrich,  1800-1882.  German  chemist.  Isolated  alu- 
minium and  beryllium.  Worked  on  boron,  silicon,  and  many  organic 
substances.  Discovered  isomerism.  Overthrew  barrier  between  or- 
ganic and  inorganic  chemistry.  Was  fellow-worker  with  Liebig,  pupil 
of  Berzelius,  and  influential  teacher  of  many  famous  chemists. 


Appendix.  447 

CHRONOLOGICAL  TABLE  OF  FAMOUS  CHEMISTS. 


Greeks 

Galen 

Aristotle 

Geber 

Avicenna 

Albertus  Magnus  Roger  Bacon 

8th  Century       978-1036 

1193-1280 

1214-1294 

Middle  Ages 

Raymond  Lulli 

Basil  Valentine 

1235-1315 

1394  

I4th  to  i6th  Cen- 

Paracelsus 

Agricola 

Libavius 

Van  Helmont 

turies. 

1493-1541 

1494-1555 

1540-1616 

1577-1644 

Glauber 

Boyle 

Becher 

Hooke 

lyth     and     i8th 

1604-1668 

1626-1691 

1635-1682 

1635-1702 

Centuries. 

Mayow 

Stahl 

Boerhaave 

Hales 

1645-1679 

1660-1734 

1668-1738 

1677-1761 

ENGLISH. 

Black 

Cavendish 

Priestley 

1728-1799 

1731-1810 

1733-1804 

Dalton 

Davy 

Faraday 

1766-1844 

1778-1829 

1791-1867 

i8th     and     igth 

FRENCH. 

Lavoisier 

Berthollet 

Proust 

Centuries. 

I743-J794 

1748-1822 

1755-1826 

Gay-Lussac 

1778-1850 

SWEDISH. 

Bergman 

Scheele 

Berzelius 

1735-1784 

1742-1786 

1779-1848 

ENGLISH. 

Graham     FRENCH.     Dumas 

BELGIAN.    Stas 

1805-1869 

1800-1884 

1813-1891 

igth  Century. 

GERMAN. 

Wohler 

Liebig 

Bunsen 

1800-1882 

1803-1873 

1811-1899 

Hofmann 

1818-1892 

5.  Atomic  Weights. — The  following  table  of  atomic  weights  is 
from  the  Journal  of  the  American  Chemical  Society,  Vol.  XXV,  No.  I 
(January,  1903). 


448 


Descriptive  Chemistry. 


TABLE  OF  ATOMIC  WEIGHTS. 


•  S 

1 

ATOMIC  WEIGHT. 

ELEMENT. 

• 

> 

o> 

O=i6. 

H«i. 

APPROXIMATE.1 

Aluminium  .... 

Al 

27.1 

26.9 

27 

Antimony 

Sb 

120.2 

"9-3 

120 

Argon  

A 

39-9 

39-6 



Arsenic         .... 

As 

75-o"~ 

74-4 

75 

Barium         .... 

Ba 

137-4 

136.4 

137 

Bismuth        .... 

Bi 

208.5 

206.9 



Boron   

B 

n. 

10.9 

11 

Bromine        .... 

Br 

79.96 

79-36 

80 

Cadmium     .... 

Cd 

112.4 

iii.6 



Caesium        .... 

Cs 

133. 

132. 



Calcium        .... 

Ca 

40.1 

39-8 

40 

Carbon          .... 

C 

I2.OO 

11.91 

12 

Cerium          .... 

Ce 

I4O. 

139. 



Chlorine        .... 

Cl 

3545 

35-18 

35.5 

Chromium    .... 

Cr 

52.1 

51-? 

52 

Cobalt  

Co 

59-o 

58.56 



Columbium  .... 

Cb 

94- 

93-3 



Copper         .... 

Cu 

63.6 

63.1 

63.5 

Erbium          .... 

Er 

166. 

164.8 



Fluorine        .        .        .      '  • 

F 

19. 

18.9 

19 

Gadolinium 

Gd 

156. 

155. 



Gallium         . 

Ga 

70. 

69-5 



Germanium 

Ge 

72.5 

71.9 



Glucinum     .... 

Gl 

9.1 

9-°3 



Gold     .        .        .        .   •     4 

Au 

197.2 

IQC.7 

197 

Helium         .... 

He 

mTM 

4- 

•*-yj'/ 
4- 

Hydrogen     .        .        .        . 

H 

1.008 

I.OOO 

1 

Indium         .... 

In 

114. 

113.1 



Iodine'  

I 

126.85 

125.00 

127 

Iridium         .... 

Ir 

193.0 

191.5 



Fe 

cqn 

rqr 

56 

Krypton        .... 

Kr 

3J';? 

81.8 

D3O 
81.2 

Lanthanum  .... 

La 

138.9 

137.9 



Lead     

Pb 

206.9 

2O^.^< 

207 

Lithium        .... 

Li 

7-03 

OOJ 

6.98 

Magnesium  .... 

Mg 

24.36 

24.18 

24 

Manganese  .... 

Mn 

55-0 

54-6 

55 

Mercury        .        .        .        . 

Hg 

200.0 

198.5 

200 

Molybdenum 

Mo 

96.0 

95-3 



1  Use  these  values  in  solving  problems. 


Appendix. 
TABLE  OE  ATOMIC  WEIGHTS  (Continued}. 


449 


• 

i 

/ 

LTOMIC  WEIGHT. 

ELEMENT. 

s 

£ 

O=i6. 

H-i. 

APPROXIMATE.1 

Neodymium 

Nd 

143.6 

142-5 



Neon    

Ne 

20. 

19.9 



Nickel 

Ni 

58.7 

58.3 



Nitrogen  

N 

14.04 

13.93 

14 

Osmium        .... 

Os 

191. 

189.6 



Oxygen         .... 

0 

16.00 

15.88 

16 

Palladium     .... 

Pd 

106.5 

105.7 



Phosphorus  .... 

P 

31.0 

30.77 

31 

Platinum       .... 

Pt 

194.8 

193-3 

195 

Potassium     .... 

K 

39-15 

38.86 

39 

Praseodymium     .        .        .  . 

Pr 

140.5 

1394 



Radium         .... 

Rd 

225. 

223.3 



Rhodium      .... 

'    Rh 

103.0 

IO2.2 



Rubidium     .        .        .    •  » 

Rb 

854 

84.8 



Ruthenium  .... 

Ru 

101.7 

100.9 



Samarium     ...» 

Sm 

150. 

148.9 



Scandium     .... 

Sc 

44.1 

43-8 



Selenium      .... 

Se 

79.2 

78.6 



Silicon  

Si 

28.4 

28.2 

28 

Silver    

Ag 

107.93 

107.12 

108 

Sodium         . 

Na 

23-05 

22.88 

23 

Strontium     ... 

Sr 

87.6 

86.94 



Sulphur         .... 

S 

32.06 

31.83 

32 

Tantalum     .... 

Ta 

183. 

181.6 



Tellurium     .... 

Te 

127.6 

126.6 



Terbium       .... 

Tb 

160. 

158.8 



Thallium      .... 

Tl 

204.1 

202.6 



Thorium       .... 

Th 

232.5 

230.8 



Thulium       .... 

Tm 

171. 

169.7 



Tin        

Sn 

119. 

118.1 

119 

Titanium      .... 

Ti 

48.1 

'47.7 



Tungsten      .... 

W 

184. 

182.6 



Uranium       .... 

U 

238-5 

236.7 



Vanadium    .... 

V 

51.2 

50.8 



Xe 

128. 

127. 



Ytterbium     .... 

Yb 

173.0 

171.7 



Yttrium         .... 

Yt 

89.0 

88.3 



Zinc      

Zn 

654- 

64.9 

65 

Zirconium    .... 

Zr 

90.6 

89.9 



1  Use  these  values  in  solving  problems. 


450  Descriptive  Chemistry. 

6.  Reference  Books  and  Supplementary  Reading.  —  The  list  of 
books  given  below  will  serve  as  the  basis  of  a  chemical  library.  The 
starred  (*)  titles  indicate  books  intended  for  the  teacher,  though  many 
parts  of  these  books  are  not  beyond  the  grasp  of  pupils.  The  library 
should  contain  at  least  numbers  i,  5,  8,  10,  18,  20,  22,  24.  Additional 
titles  can  be  found  in  (i)  List  of  Books  in  Chemistry,  L.  E.  Knott 
Apparatus  Co.,  Boston,  Mass. ;  (2)  Smith  and  Hall's  Teaching  of 
Chemistry  and  Physics,  p.  218;  (3)  NEWELL'S  EXPERIMENTAL  CHEM- 
ISTRY, APP.  C,  II. 

i.  Text-Book  of  Inorganic  Chemistry,  Newth.  Longmans,  Green, 
&  Co.,  682  pp.,  $1.75. 

*2.  General  Inorganic  Chemistry,  Freer.  Allyn  &  Bacon,  Boston, 
559  PP-»  #3- 

*3.  Text-Book  of  Inorganic  Chemistry,  Holleman.  John  Wiley  & 
Sons,  458  pp.,  $2.50. 

4.  Physical  Chemistry  for  Beginners,  Van  Deventer.     John  Wiley 
&  Sons,  154  pp.,  $1.50. 

5.  Chemical  Theory  for  Beginners,  Dobbin  and  Walker.      The 
Macmillan  Co.,  236  pp.,  $  .70. 

*6.  Introduction  to  Physical  Chemistry,  Walker.  The  Macmillan 
Co.,  332  pp.,  $3. 

7.  The  Birth  of  Chemistry,  Rodwell.      The  Macmillan  Co.,  135 
pp.,  $i. 

8.  Short  History  of  Chemistry,  Venable.     D.  C.  Heath  &  Co.,  172 
pp.,  $i. 

9.  Faraday  as  a  Discoverer,  Tyndall.     D.  Appleton  &  Co.,  171 
pp.,  $i. 

10.  Short  History  of  Natural  Science,  Buckley.     D.  Appleton  & 

CO.,    467    pp.,    $2. 

11.  Heroes  of  Science  —  Chemists,  Muir.     Thomas  Nelson  &  Son, 
350  pp.,  $1.50. 

*I2.  Essays  in  Historical  Chemistry,  Thorpe.  The  Macmillan  Co., 
582  pp.,  $4. 

13.  Humphry  Davy,  Thorpe.     The  Macmillan  Co.,  240  pp.,  $1.25. 

14.  John  Dalton,  Roscoe.     The  Macmillan  Co.,  216  pp.,  $1.25. 

15.  Michael   Faraday,  Thompson.     The  Macmillan  Co.,  308  pp., 
$1.25. 

*i6.  Alembic  Club  Reprints,  University  of  Chicago  Press,  $.40  each, 
(i)  Experiments  on  Magnesia  Alba.  (2)  Foundations  of  the  Atomic 


Appendix.  451 

Theory.  (3)  Experiments  on  Air.  (4)  Foundations  of  the  Molecular 
Theory.  (6)  Decomposition  of  the  Fixed  Alkalies.  (7)  (8)  Discov- 
ery of  Oxygen.  (9)  Elementary  Nature  of  Chlorine.  (13)  Early  His- 
tory of  Chlorine. 

*ij.  Organic  Chemistry,  Remsen.  D.  C.  Heath  &  Co.,  426  pp., 
$1.30. 

18.    Outlines  of  Industrial  Chemistry,  F.  H.  Thorp.     The  Macmil- 
lan  Co.,  528  pp.,  $3.50. 

*I9-  Practical  Electro-Chemistry,  Blount.  The  Macmillan  Co., 
374  pp.,  $3.25. 

20.  Chemistry  in  Daily  Life,  Lassar-Cohn.     J.  B.  Lippincott  Co., 
336pp.,  $1.75. 

21.  The  Soil,  King.     The  Macmillan  Co.,  400  pp. 

22.  Story  of  a  Piece  of  Coal,  Martin.     D.  Appleton  &  Co.,  165  pp., 
$.40. 

23.  Chemical  History  of  a  Candle,  Faraday.     Harper  &  Bros.,  223 
pp.,  $1.00. 

24.  Minerals  and  How  to  Study  Them,  E.  S.  Dana.    John  Wiley  & 
Sons,  380  pp.,  $1.25. 

*25.  Teaching  of  Chemistry  and  Physics,  Smith  and  Hall.  Long- 
mans, Green  &  Co.,  384  pp.,  $1.50. 

26.  Story   of   Nineteenth-Century  Science,  Williams.     Harper  & 
Bros.,  475  PP->  $2.50. 

27.  Stories  of  Industry,  Vol.  I,  Chase  and  Clow.     Educational  Pub- 
lishing Co.,  Boston,  172  pp.,  $.40. 

Scientific  American,  Munn  &  Co.,  New  York.  $3.00  yearly;  single 
copies,  8  cents. 

School  Science,  Ravenswood,  Chicago,  Illinois.  $2.00  yearly  (9 
issues)  ;  single  copies,  25  cents. 

Popular  Science  Monthly,  The  Science  Press,  New  York.  $3.00 
yearly ;  single  copies,  25  cents. 


PART   II 
EXPERIMENTS 


CONTENTS. 

PART   II. 

(Numbers  in  parentheses  indicate  experiments.) 

PAGE 

INTRODUCTION        .        .        .        .        .        .        .....        .    459 

Bunsen  Burner  ;  Heating  ;  Cutting  and  Bending  Glass  Tubing ; 
Filtering  ;  Constructing  and  Arranging  Apparatus  ;  Manipula- 
tion ;  Smelling  and  Tasting. 

PHYSICAL  AND  CHEMICAL  CHANGES      .        .    "    .   '    .        .        .        .    467 

Physical  Change  (i,  2,  3);   Chemical  Change  (4). 

OXYGEN  .        .        ...        .        .        .        ...        .        .    468 

Preparation  (5) ;   Properties  (6) ;   Preparation   from  Mercuric  Oxide 

(7)- 

HYDROGEN     .        .        .       -.        .        .        .        .        .        .       ;.        .    471 

Preparation  (8);   Properties  (9);   Burning  Hydrogen  (10). 

WATER   .        .  .        .        .        .        .        .        .        .        .    474 

General  Distribution  (n);  Tests  for  Impurities  (12);  Distillation 
(13);  Solubility  of  Gases  (14);  Solubility  of  Liquids  (15);  Solu- 
bility of  Solids  (16);  Supersaturation  (17);  Water  of  Crystal- 
lization (18);  Efflorescence  (19);  Deliquescence  (20);  Solution 
and  Chemical  Action  (21) ;  Electrolysis  (22) ;  Water  and  Chloripe 
(23);  Water  and  Sodium  (24). 

THE  AIR ,  481 

Composition  (25);   Water  Vapor  (26) ;   Carbon  Dioxide  (27). 

ACIDS,  BASES,  AND  SALTS 483 

Properties  of  Acids  (28) ;  Properties  of  Bases  (29)  ;  A  Property  of 
Salts  (30);  Nature  of  Common  Substances  (31);  Neutralization 
(32). 

HEAT,  LIGHT,  ELECTRICITY,  AND  CHEMICAL  ACTION   .....    485 
Heat  and  Chemical  Action   (33,  34) ;    Light  and  Chemical  Action 
(35) ;   Electricity  and  Chemical  Action  (36,  37). 

455 


456  Descriptive  Chemistry. 

PAGE 

CHLORINE 486 

Preparation  (38) ;  Properties  (39) ;  Bleaching  Powder  (40) ;  Prep- 
aration of  Hydrochloric  Acid  (41);  Properties  of  Hydrochloric 
Acid  Gas  (42)  ;  Properties  of  Hydrochloric  Acid  (43);  Tests  for 
Hydrochloric  Acid  and  Chlorides  (44). 

COMPOUNDS  OF  NITROGEN .'       .        .        .    490 

Preparation  of  Ammonia  (45);  Properties  of  Ammonia  Gas  (46); 
Properties  of  Ammonium  Hydroxide  (47) ;  Neutralization  of  Am- 
monia (48) ;  Preparation  of  Nitric  Acid  (49) ;  Properties  of  Nitric 
Acid  (50);  Test  for  Nitric  Acid  and  Nitrates  (51-52);  Interaction 
of  Sodium  Nitrate  and  Sulphuric  Acid  (53) ;  Nitric  Acid  and  Metals 
(54);  Nitric  Acid  and  Copper,  and  Nitrogen  Peroxide  (55);  Ni- 
trous Oxide  (56);  Sodium  Nitrite  (57);  Aqua  Regia  (58). 

CARBON .,...'.        .        .    498 

Distribution  (59) ;  Decolorizing  Action  (60) ;  Deodorizing  Action 
(61 ) ;  Preparation  of  Carbon  Dioxide  (62)  ;  Properties  of  Carbon 
Dioxide  (63) ;  Interaction  of  Calcium  Carbonate  and  Hydro- 
chloric Acid  (64);  Carbon  Dioxide  and  Combustion  (65);  Car- 
bonic Acid  (66)  ;  Carbonates  (67)  ;  Detection  of  Carbonates  (68) ; 
Acid  Calcium  Carbonate  (69);  Carbon  Monoxide  (70);  Ethylene 
(71);  Acetylene  (72);  Illuminating  Gas  (73);  Combustion  of 
Illuminating  Gas  (74);  Bunsen  Burner  (75);  Bunsen  Burner 
Flame  (76);  Candle  Flame  (77);  Kindling  Temperature  (78); 
Reduction  and  Oxidation  (79). 

FLUORINE,  BROMINE,  AND  IODINE        . 511 

Hydrofluoric  Acid  (80);  Bromine  (81);  Potassium  Bromide  (82); 
Iodine  (83);  Tests  for  Iodine  (84,  85);  Detection  of  Starch 
(86);  Potassium  Iodide  (87). 

SULPHUR        .        . '•••>   .        .       .       .       ...        .        .        .    514 

Properties  (88) ;  Amorphous  Sulphur  (89) ;  Crystallized  Sulphur  (90) ; 
Combining  Power  (91);  Sulphur  and  Matches  (92);  Preparation 
of  Hydrogen  Sulphide  (93);  Properties  of  Hydrogen  Sulphide 
Gas  (94)  ;  Sulphides  (95) ;  Preparation  of  Sulphur  Dioxide  (96) ; 
Properties  of  Sulphur  Dioxide  Gas  (97) ;  Properties  of  Sulphurous 
Acid  (98);  Sulphuric  Acid  and  Organic  Matter  (99);  Test  for 
Sulphuric  Acid  and  Sulphates  (100). 


Contents.  457 

PAGE 

SILICON  AND  BORON 520 

Silicic  Acid  (101);  Borax  Beads  (102);  Boric  Acid  (103). 

PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH 522 

Properties  of  Phosphorus  (104);  Test  for  Arsenic  (105);  Test  for 
Antimony  (106);  Test  for  Bismuth  (107). 

SODIUM  AND  POTASSIUM         . 522 

Properties  of  Sodium  (108);  Sodium  Hydroxide  (109);  Exercises; 
Properties  of  Potassium  (no);  Potassium  Hydroxide  (in);  Potas- 
sium Carbonate  (112);  Exercises. 

COPPER,  SILVER,  AND  GOLD  ,  •••    .     •   , , 525 

Properties  of  Copper  (113);  Tests  for  Copper  (114);  Interaction  of 
Copper  with  Metals  (115);  Exercises;  Preparation  of  Silver  (116); 
Properties  of  Silver  (117);  Test  for  Silver  (118);  Exercises  ;  Test 
for  Gold  (119). 

CALCIUM,  STRONTIUM,  AND  BARIUM      .        ....        .        .        .     528 

Tests  for  Calcium  ( 1 20) ;  Plaster  of  Paris  (121) ;  Exercises;  Test  for 
Strontium  (122);  Red  Fire  (123);  Tests  for  Barium  (124);  Green 
Fire  (125);  Exercises. 

MAGNESIUM,  ZINC,  CADMIUM,  AND  MERCURY  .  .  .  /  .  530 
Properties  of  *  Magnesium  (126);  Tests  for  Magnesium  (127); 
Exercises  ;  Properties  of  Zinc  (128);  Tests  for  Zinc  (129);  Inter- 
action of  Zinc  and  Metals  (130);  Exercises;  Test  for  Cadmium 
(131);  Properties  of  Mercury  (132);  Tests  for  Mercury  (133); 
Mercurous  and  Mercuric  Compounds  (134);  Exercises. 

ALUMINIUM .        .        .     532 

Properties  (135);  Action  with  Acids  and  Alkalies  (136);  Aluminium 
Hydroxide  (137);  Tests  (138);  Alum  (139). 

TIN  AND  LEAD 534 

Properties  of  Tin  (140);  Action  of  Tin  with  Acids  (141);  Tests  for 
Tin  (142);  Deposition  (143);  Properties  of  Lead  (144);  Tests 
for  Lead  (145);  Deposition  of  Lead  (146);  Oxides  of  Lead  (147); 
Compounds  of  Lead  (148). 

CHROMIUM  AND  MANGANESE 537 

Tests  for  Chromium  (149);  Chromates  (150);  Reduction  of  Chro- 
mates  (151);  Chromic  Hydroxide  (152);  Chrome  Alum  (153); 
Tests  for  Manganese  (154);  Potassium  Permanganate  (155); 
Exercises. 


458  Descriptive  Chemistry. 

PAGE 

IRON,  NICKEL,  AND  COBALT .     540 

Properties  of  Iron  (156);  Ferrous  Compounds  (157);  Ferric  Com- 
pounds (158);  Reduction  of  Ferric  Compounds  (159);  Oxidation 
of  Ferrous  Compounds  ( 1 60);  Compounds  of  Iron  (161);  Exercises; 
Test  for  Nickel  (162);  Test  for  Cobalt  (163). 

ORGANIC  COMPOUNDS .        .    542 

Composition  (164);  Alcohol  (165);  Properties  of  Alcohol  (166); 
Aldehydes  (167);  Ether  (168);  Acetic  Acid  (169);  Vinegar 
(170);  Test  for  Acetic  Acid  and  Acetates  (171);  Acetates  (172); 
Organic  Acids  (173);  Ethyl  Acetate  (174);  Soap  (175,  176); 
Glycerine  (177);  Test  for  Sugar  (178);  Exercises  ;  Benzene  (179). 

LABORATORY  EQUIPMENT       .        .        ...       .       .        .        .    549 

Apparatus  ;   Chemicals  ;   Solutions. 


INTRODUCTION. 


1.  The  Bunsen  burner  is  used  as  the  source  of  heat  in  most  chem- 
ical laboratories  (Fig.  87).  It  is  attached  to  the  gas  cock  by  a  piece  of 
rubber  tubing.  When  the  gas  is  turned  on,  the  current  of  gas  draws 
air  through  the  holes  at  the  bottom  of  the  tube,  and  this  mixture  when 
lighted  burns  with  an  almost  colorless,  /.  e.  non-luminous,  flame.  It  is  a 
hot  flame  and  deposits  no  soot.  The  burner  is  lighted  by  turning  on 
the  gas  full  and  holding  a  lighted  match  in  the  gas  about  5  centimeters 
(2  inches)  above  the  top  of  the  burner.  If  the 
flame  is  not  colorless,  or  nearly  so,  turn  the  ring  at 
the  bottom  of  the  burner  until  the  flame  is  a  faint 
blue.  The  colorless  flame  should  be  used  in  all 
experiments  unless  the  directions  state  otherwise, 
and  should  be  from  5  to  10  centimeters  (2  to  4 
inches)  high.  The  hottest  part  of  the  flame  is  near 
the  top. 


FIG.  87.  —  Bunsen 
burner. 


2.    Heating. — The  following  directions  should 
be  observed  in  heating  with  the  Bunsen  burner  :  — 

(1)  The  burner  should  always  be  lighted  before 

any  piece  of  apparatus  is  held  over  it,  or  before  it  is  placed  beneath 
a  wire  gauze  which  supports  a  dish  or  flask. 

(2)  Glass  and  porcelain  apparatus  should  not  be  heated  when  empty 
nor  over  a  bare  or  free  flame  even  if  they  contain  something  —  unless 
directions  so  state.     Vessels  requiring  a  support  should  be  placed  on  a 
wire  gauze  which  stands  on  the  ring  of  an  iron  stand,  and  heated  grad- 
ually from  beneath.     Hot  vessels  should  be  heated  and  cooled  gradu- 
ally ;  if  removed  from  the  gauze  while  hot,  they  should  be  placed  on  a 
block  of  wood  or  piece  of  asbestos  board  —  never  on  a  cold  surface. 

(3)  Many  experiments  require  the   heating  of  test   tubes.     These 
tubes  should  be  dry  on  the  outside  before  being  heated.     The  temper- 
ature of  a  test  tube  containing  a  solid  should  be  raised  gradually  by 
moving  it  in  and  out  of  the  flame,  or  by  holding  it  in  the  flame  and  roll- 

459 


460 


Experiments. 


ing  it  slightly  between  the  thumb  and  forefinger.  Special  care  must  be 
taken  to  distribute  the  heat  evenly.  If  the  test  tube  contains  a  liquid, 
as  is  usually  the  case,  only  that  part  containing  the  liquid  should  be 
heated ;  the  test  tube  should  also  be  inclined  so  that  the  greatest  heat 
is  not  directed  upon  the  thin  bottom. 
When  the  liquid  begins  to  boil,  the  test 
tube  should  be  removed  from  the  flame 
for  an  instant  or  held  over  it.  In  some 
experiments  test  tubes  can  be  held  be- 
tween the  thumb  and  forefinger  without 
discomfort.  If  they  are  too  hot  to 
handle,  a  test-tube  holder  may  be  used  (Fig.  88). 

3.  Cutting  and  bending  Glass  Tubing.  —  (a)  Cut- 
ting. Determine  the  length  needed,  lay  the  tube  on  the 
desk,  and  with  a  forward  stroke  of  a  triangular  file  make 
a  short  but  deep  scratch  where  the  tube  is  to  be  cut. 
Grasp  the  tube  in  both  hands,  and  hold  the  thumbs 
together  behind  the  scratch.  Now  push  gently  with  the 
thumbs,  pull  at  the  same  time  with  the  hands,  and  the 
tube  will  break  at  the  desired  point.  The  sharp  ends 
should  be  smoothed  by  rubbing  them  with  emery  paper 
or  by  rotating  them  slowly  in  the  Bunsen  flame  until  a 
yellow  color  is  distinctly  seen  or  until  the  ends  become  red-hot. 

(£)  Bending.  Glass  tubes  are  bent  in  a  flat  flame. 
An  ordinary  illuminating  gas  flame  may  be  used, 
but  the  Bunsen  flame  can  be  flattened  by  a  wing-top 
attachment  (Fig.  89),  which  slips  over  the  top  of  the 
burner  tube. 
The  flattened 
Bunsen  flame 

should  be  slightly  yellow  and 
about  7  centimeters  (2.5  inches) 
wide  for  ordinary  bends.  A 
right-angle  bend  is  made  as 
follows:  Determine  the  point 
at  which  the  tube  is  to  be  bent. 
Grasp  the  tube  in  both  hands, 
and  hold  it  so  that  the  part  to 


FIG.  88.— 
Test  tube  and 
holder. 


FIG.  89.  —  Wing- 
top  attachment  for 
Bunsen  burner. 


be  bent  is   directly   over  the 


FIG.  90.  — Bending  a  tube  into  a  right 
angle  — I. 


Introduction.  461 

flame.  Slowly  rotate  it  between  the  thumbs  and  forefingers,  and 
gradually  lower  it  into  the  position  shown  in  Figure  90.  Continue  to 
rotate  it  until  the  glass  feels  soft  and  ready  to  yield.  Then  remove 
it  from  the  flame,  and  slowly 
bend  it  into  a  right  angle,  as 
shown  in  Figure  91.  It  is  con- 
venient to  have  at  hand  a  block 
of  wood  or  some  other  right- 
angled  object  to  assist  the  eye 

in  completing  the  bend  into  an 

.    ,.          ,         Tr      „  FIG.  91.  —  Bending  a  tube  into  a  right 

exact  right  angle.     If  a  Bunsen  angle  —  II. 

flame  is  used,  the  bent  part  of 

the  tube  should  be  annealed,  i.e.  cooled  slowly.  This  is  done  by 
holding  it  in  a  yellow  flame  until  it  becomes  coated  with  soot.  It 
should  then  be  placed  on  a  block  of  wood,  and  when  cold  wiped 
clean.  Tubes  can  be  bent  into  an  oblique  angle  by  heating  them 
through  about  twice  the  space  required  for  a  right  angle ;  a  very  slight 
bend,  however,  is  often  made  by  holding  the  tube  across  the  flame  and 
heating  a  short  space.  Glass  tubes  which  have  been  correctly  bent 
never  have  flattened  curves ;  nor  are  they  twisted,  i.e.  all  parts  lie  in 
the  same  plane. 

(c)  Drawing.  Glass  tubes  can  be  drawn  to  a  finer  bore  or  into  two 
pointed  tubes  as  follows :  Heat  the  glass  as  in  (b)  through  about 
2.5  centimeters  (i  inch)  of  its  length,  remove  from  the  flame  and 
slowly  pull  it  apart  a  short  distance ;  let  it  cool  for  a  few  seconds,  and 
then  pull  it  quickly  to  the  desired  length. 

The  operation  is  well  illustrated  by  making  a  glass  stirring  rod. 
Select  a  piece  of  rod  about  25  centimeters  (10  inches)  long  and  .5 


FIG.  92.  —  Stirring  rods  ready  to  be  cut. 

centimeter  (^  inch)  in  diameter.  Heat  it  in  the  middle  in  the 
ordinary  —  not  flat  —  Bunsen  flame,  and  when  soft  draw  it  out  slowly 
into  the  shape  shown  in  Figure  92.  Cut  it  into  two  rods  by  making  a 
slight  scratch  where  the  dotted  line  indicates.  Round  off  the  rough 
edges  by  heating  them  slightly  in  the  flame. 


462 


Experiments. 


4.  Filtering.  —  A  solid  may  be  separated  from  a  liquid  by  filtering. 
A  circular  piece  of  porous  paper  is  folded  to  fit  a  glass  funnel,  and  when 
the  mixture  is  poured  upon  this  paper,  the  solid — the  residue  or  precipi- 
tate—  is  retained,  while  the  liquid — the  filtrate  —  passes  through  and 
may  be  caught  in  a  test  tube  or  any  other  vessel.  The  filter  paper  is 
prepared  for  the  funnel  by  folding  it  successively  into  the  shapes  shown 
in  Figures  93  and  94,  and  then  opening  the  folded  paper  so  that  three 
thicknesses  are  on  one  side  and  one  on  the  other  (Fig.  95).  The 
cone-shaped  paper  is  next  placed  in  the  funnel  and  wet  with  water, 


FIG.  93.  —  Folded 
filter  paper  —  I. 


FIG.  94.  —  Folded 
filter  paper — II. 


FIG.  95.  — Folded  filter 
paper  ready  for  funnel. 


so  that  it  will  stick  to  the  sides  of  the  funnel  and  filter  rapidly.  The 
paper  should  never  extend  above  the  edges  of  the  funnel,  but  its  apex 
should  always  project  slightly  into  the  stem.  The  liquid  to  be  filtered 
should  be  poured  down  a  glass  rod  which  touches  the  edge  of  the  test 
tube ;  the  lower  end  of  the  rod  should  just  touch  the  paper  inside  the 
funnel,  so  that  the  liquid  will  run  down  the  side  and  thereby  avoid 
bursting  the  apex  of  the  filter  paper.  It  is  also  advisable  to  adjust  the 
apparatus  so  that  the  end  of  the  stem  of  the  funnel  rests  against  the 
side  of  the  vessel  catching  the  filtrate.  A  funnel  can  be  supported  by 
standing  it  in  a  test  tube,  a  bottle,  or  the  ring  of  an  iron  stand. 

5.  Constructing  and  arranging  Apparatus.  —  The  various  parts 
of  an  apparatus  should  be  collected,  prepared,  and  put  together  be- 
fore starting  the  experiment  in  which  the  apparatus  as  a  whole  is 
used.  The  different  parts  which  are  to  fit  each  other  should  be  selected 
and  arranged  so  that  all  joints  are  gas-tight,  and  as  a  final  precaution 
the  apparatus  should  be  tested  for  leaks.  All  leaks  should  be  stopped 
up  before  the  apparatus  is  used.  The  following  hints  will  be  helpful :  — 


FIG.  96.  —  Rubber  tube  cut  at  an  angle. 

(1)    To  insert  a  glass  tube  into  rubber  tubing.     Cut  the  rubber  tubing 
at  an  angle,  as  shown  in  Figure  96,  moisten  the  smoothed  end  of  the  glass 


Introduction.  463 

tube  with  water,  place  the  end  of  the  glass  tube  in  the  angular-shaped 
cavity  so  that  both  tubes  are  at  about  a  right  angle,  and  then  slip  the 
rubber  tube  slowly  up  and  over  the  end  of  the  glass  tube.  If  the  glass 
tube  is  large  or  the  rubber  stiff,  the  rubber  tube  must  be  held  firmly 
between  the  thumb  and  forefinger  to  keep  it  from  slipping  off  until  it  is 
securely  adjusted. 

(2)  To  fit  a  glass  tube  to  a  stopper.     Moisten  the  end  with  water  and 
grasp  the  tube  firmly  about  3  centimeters  (i   inch)  from  the  end;  hold 
the  stopper  between  the  thumb  and 'forefinger  of  the  other  hand,  and 
work  the  tube  into  the  hole  by  a  gradual  rotary  motion.     Proceed  in 
the  same  manner  if  the  tube    is  to  be  pushed  through  the  stopper. 
Never  point  the  tube  toward  the  palm  of  the  hand  which  holds  the 
stopper.     Never  grasp  a  safety  tube  or  any  bent  tube  at  the  bend  when 
inserting  it  into  a  stopper  —  it  may  break  and  cut  the  hand  severely. 

(3)  To  bore  a  hole  in  a  cork.     Rubber  stoppers  are  preferable,  but 
if  corks  are  used,  they  can  be  bored  as  follows :  Select  a  cork  free  from 
cracks  or  channels  and  use  a  borer  which  is  one  size  smaller  than  the 
desired  hole.     Hold  the  cork  between  the  thumb  and  forefinger,  press 
the  larger  end  against  a  firm  but  soft  board,  and  slowly  push  the  borer 
by  a  rotary  movement  through  the  cork,  taking  care  to  keep  the  borer 
perpendicular  to  the  cork.     If  the  hole  is  too  small,  enlarge  it  with  a 
round  file.     If  corks  are  used  instead  of  rubber  stoppers,  the  apparatus 
should  always  be  tested  before  use  by  blowing  into  it,  stopping  of  course 
all  legitimate  outlets.     A  poor  cork  often  means  a  failure,  to  say  noth- 
ing of  wasted  time. 

(4)  To  make  a  platinum  test  wire.     Rotate  one  end  of  a  piece  of 
glass  rod,  about  10  centimeters  (4  inches)  long,  in  the  flame  until  it 
softens.     At  the  same  time  grasp  a  piece  of  platinum  wire  about  7  cen- 
timeters (3  inches)  long  firmly  in  the  forceps  about  i  centimeter  (.5 
inch)  from  the  end,  and  hold  it  in  the  flame.     When  the  rod  is  soft 
enough,  gently  push  the  hot  wire  into  the  rod.     Cool  the  rod  gradually 


FIG.  97.  —  Platinum  test  wire. 

by  rotating  it  in  the  flame.  The  completed  wire  is  shown  in  Figure  97. 
If  a  glass  tube  is  used  instead  of  a  rod,  it  should  be  drawn  out  to  a 
very  small  diameter  (see  §  3  (0)  before  inserting  the  platinum  wire, 
but  in  other  respects  the  two  operations  are  practically  identical. 


464 


Experiments. 


6.  Manipulation.  —  Ability  to  use  apparatus  rapidly,  accurately,  and 
neatly  is  acquired  only  by  experience,  but  the  following  suggestions  will 
facilitate  the  acquisition  of  this  needful  skill :  — 

(i)  Pouring  liquids  and  transferring  solids,     (a)    Liquids  can  be 

poured  from  a  vessel  without 
spilling,  by  moistening  a  glass 
rod  with  the  liquid  and  then 
pouring  it  down  the  rod  as 
is  shown  in  Figure  98.  The 
angle  at  which  the  rod  is  held 
varies  with  circumstances. 

This  is  a  convenient  way  to 
FIG.  98. —  Pouring  a  liquid  down  a  glass  rod.  i-      -j   r  i 

pour   a  liquid  from  a   vessel 

containing  a  solid  without  disturbing  the  solid.  (£)  Liquids  can  often 
be  poured  from  a  bottle  by  holding  the  bottle  as  shown  in  Figure  99. 
Notice  that  the  stopper  and  bottle  are  held  in  the  same  hand.  This  is  ac- 
complished by  holding  the 
palm  of  the  hand  upward 
and  removing  the  stopper 
by  grasping  it  between  the 
fingers  before  the  bottle  is 
lifted.  All  stoppers  should 
be  removed  this  way  when 
possible,  and  not  laid  down, 
because  the  impurities  ad- 
hering to  the  stopper  may 
run  down  into  the  bottle 
and  contaminate  the  solu- 
tion. The  drop  on  the  lip  of  the  bottle  should  be  touched  with  the 
stopper  before  the  latter  is  put  into  the  bottle ;  this  simple  operation 

prevents  the  drop  from  running 
down  the  outside  of  the  bottle 
upon  the  label  or  upon  the 
shelf.  (<:)  Solids  should  never 
be  poured  directly  from  a  large 
bottle  into  a  test  tube,  retort,  or 
similar  vessel.  A  convenient 
method  is  as  follows :  Rotate 


FIG.  99.  —  The  way  in  which  a  glass  stopper 
should  be  held  while  a  liquid  is  being  poured 
from  a  bottle. 


FlG.  loo.  —  Pouring  a  solid  into  a  vessel  with 
a  small  opening. 


the  bottle  slowly  so  that  the 


Introduction. 


465 


solid  will  roll  out  in  small  quantities ;  catch  the  solid  on  a  narrow  strip 
of  paper  folded  lengthwise,  and  slide  the  solid  from  the  paper  into  the 
desired  vessel.  The  last  part  of  the  operation  is  shown  in  Figure  100. 

(2)  Collecting  gases.     Gases   are   usually   collected   over  water    by 
means  of  a  pneumatic  trough,  a  common  form  of  which  is  shown  in  Figure 
102.     The  vessel  to  be  filled  with  gas  is  first  filled  with  water,  covered 
with  a  piece  of  filter  paper,  inverted,  and  placed  mouth  downward  on  the 
shelf  of  the  trough,  which  is  previously  filled  with  water  just  above  the 
shelf.     The  paper  is  then  removed,  and  the  vessel  slipped  over  the  hole 
in  the  shelf  of  the  trough.     Glass  plates  instead  of  filter  paper  may  be 
used  to  cover  the  bottle.     The  gas  which  is  evolved  in  the  generator 
passes  through  the  delivery  tube,  and  bubbles  up  through  the  water  into 
the  bottle,  forcing  the  water  out  of  the  bottle  as  it  rises.     All  gases 
insoluble   in  water   are   thus  collected.       Some  heavy  gases,  such  as 
hydrochloric  acid,  chlorine,  and  sulphur  dioxide,  are  collected  by  allowing 
the  gas  to  flow  downward  into  an   empty  bottle,  and  displace  the  air  in 
the  bottle,  i.e.  by  downward  displacement.     Ammonia  and  other  light 
gases  are  usually  collected  by  allowing  the  gas  to  flow  upward  into  a 
bottle,  i.e.  by  upward  displacement. 

(3)  Weighing  and  measuring.     These  operations  are  best  learned 
by  personal  direction  from  the  teacher,  together  with  patient  application 
of  a  few  general  principles.     The  following  hints,  however,  will  be  of 
assistance :  — 

(a)  Learn  as  soon  as  possible  how  to  use  the  scales  and  interpret  the 
weights. 

(b)  Always  leave  the  scales  and  weights  in  a  clean,  usable  condition. 

(c)  Substances  should  not  be  weighed 
on  the  bare  scale  pan,  but  on  a  smooth 
piece   of  paper  creased  on  the  edges  or 
along  the  middle.     Take  the  solid  from 

the  bottle  with  a  clean  spoon  or  spatula  or     --- 

pour  by  rotating  the  bottle  as  described  in 

§  6  (c).     In  many   experiments  only  ap-       -"          12- 

proximate  quantities  are  needed.     If  you 

weigh  out  too  much,  do  not  put  it  back 

into  the  bottle,  but  throw  it  away  or  put  it 

into  a  special  bottle. 

(d)  Liquids  are  measured  in  graduated     FlG.  IOI.^  Meniscus.   Correct 
cylinders.    The  lowest  point  of  the  curved  reading  is  along  line  I. 


466  Experiments. 

surface  of  the  liquid  is  its  correct  height  (see  Fig.  101).  The  average 
ordinary  test  tube  holds  about  30  cubic  centimeters,  while  the  large  test 
tube — so  often  mentioned  in  the  succeeding  experiments  —  holds  about 
75  cubic  centimeters.  Time  can  be  saved  by  remembering  these  volumes. 
(>)  All  measurements  in  this  book  are  in  the  metric  system  (see  App. 
§  i).  The  common  denominations,  their  abbreviations,  and  English 
equivalents  should  be  learned. 

7.  Smelling  and  Tasting.  —  Unfamiliar  substances  should  never 
be  tasted  or  smelled  except  according  to  directions,  and  even  then 
with  the  utmost  caution.  Never  inhale  a  gas  vigorously,  but  waft  it 
gently  with  the  hand  toward  the  nose.  Taste  acids,  etc.,  by  touching  a 
minute  portion  to  the  tip  of  the  tongue,  and  as  soon  as  the  sensation  is 
detected,  reject  the  solution  at  once  —  never  swallow  it. 


EXPERIMENTS. 

PHYSICAL  AND   CHEMICAL  CHANGES. 

Experiment  1.  — Physical  Change.    Materials:  Sugar,  glass  rod. 

Dissolve  a  little  sugar  in  a  test  tube  one  fourth  full  of  water.  Dip  a 
glass  rod  into  the  liquid  and  taste  it.  Has  the  characteristic  property 
of  the  sugar  been  changed  ?  Dip  the  rod  into  the  liquid  again,  and 
hold  it  over  the  flame  of  the  Bunsen  burner.  'As  the  water  evaporates, 
a  white  solid  appears.  Taste  it.  What  is  it  ?  Have  its  original  proper- 
ties been  destroyed  ?  What  kind  of  a  change  did  they  undergo  ? 
What  kind  of  a  change  did  the  sugar  undergo  ?  What  caused  the 
change  ? 

Experiment  2.  —  Physical  Change.     Material:  Iodine. 

Drop  a  small  crystal  of  iodine  into  a  dry  test  tube,  and  gently  heat 
the  bottom.  As  the  violet  vapor  arises,  remove  the  tube  from  the  flame 
and  let  it  cool.  Do  the  crystals  which  form  near  the  top  resemble  the 
original  crystal  ?  When  gently  heated,  do  they  change  into  the  violet 
vapor  ?  How  has  the  iodine  crystal  been  changed  ?  What  caused  the 
change  ?  Do  the  original  properties  reappear  after  the  cause  has  been 
removed  ?  What  kind  of  a  change  has  the  iodine  undergone  ? 

Experiment  3.  —  Physical  Change.     Material:  Glass  rod. 

Rub  a  glass  rod  briskly  on  a  piece  of  cloth,  and  hold  it  near  small 
bits  of  dry  paper.  Describe  what  happens.  After  a  moment  touch  the 
paper  again.  Is  the  result  the  same  ?  Try  again.  Are  the  original 
properties  of  the  glass  restored  when  the  cause  of  its  change  is  re- 
moved ?  What  kind  of  a  change  did  the  glass  undergo  ? 

Experiment  4.  —  Chemical  Change.  Materials :  Copper  wire,  dilute 
nitric  acid,  iron  nail,  forceps. 

(a)  Examine  a  piece  of  copper  wire  and  notice  especially  its  color. 
Grasp  one  end  of  the  wire  with  the  forceps,  and  hold  the  other  end  in 
1  467 


468  Experiments. 

the  flame  until  a  definite  change  occurs.  Then  remove  it  from  the 
flame,  and  examine.  Has  it  been  changed  ?  Do  the  original  properties 
of  the  copper  reappear  when  the  heated  wire  is  cool  ?  What  kind  of  a 
change  has  the  copper  undergone  ?  Has  the  change  produced  another 
substance  ? 

(^)  Slip  another  piece  of  copper  wire  into  a  test  tube  one  fourth  full 
of  dilute  nitric  acid.  Notice  any  change.  Warm  the  liquid  gently, 
and  notice  any  additional  change.  What  are  the  evidences  of  chemical 
change  ?  What  caused  the  change  ?  What  assisted  or  hastened  it  ? 
How  has  the  copper  been  changed  ?  (Save  the  test  tube  and  contents 
for  (,).) 

(V)  Carefully  slip  an  iron  nail  into  the  liquid  remaining  from  ($)  ;  let 
it  stand  a  short  time.  Then  remove  and  examine  the  coating.  How 
does  it  compare  with  the  original  copper  used  in  (a)  ?  What  kind  of 
a  change  occurred  ?  What  caused  it  ? 

ANSWER  : 

(1)  What  are  the  evidences  of  chemical  changes  in  this  experiment? 

(2)  If  a  known  weight  of  copper  had  been  consumed  in  (£),  could  it 
have  been  obtained  without  loss  in  (c)  ? 

(3)  Did  the  changes  in  this  experiment  involve  any  loss  of  copper? 

(4)  What  is  the  evidence  that  new  substances  were  produced  in  (#) 
and  (£)  ? 

(5)  What  physical  changes  occurred  in  (#)  and  (fr)  ? 

OXYGEN. 

Experiment^.  —  Preparation  of  Oxygen.  Materials:  jjjjgrarns 
potassium  chlorate,  1 5  grams  manganese  dioxide,  g^bottles  (about  250 
cubic  centimeters  each),  filter  paper,  thin  piece  of  soft  wood,  sulphur, 
deflagrating  spoon,  piece  of  charcoal  fastened  to  a  wire,  piece  (about  1 5 
centimeters  or  6  inches)  of  wire  picture  cord  unwound  at  one  end.  The 
apparatus  is  shown  in  Figure  102.  A  is  a  large  test  tube  provided  with 
a  one-hole  rubber  stopper,  to  which  is  fitted  a  short  glass  tube,  B ;  the 
delivery  tube.  I),  is  attached  to  the  short  glass  tube  by  the  rubber 
tube,  C.  (Directions  for  constructing  and  arranging  the  apparatus  may 
be  found  in  the  Introduction,  §  5.) 

Weigh  the  potassium  chlorate  on  a  piece  of  paper  creased  lengthwise, 
and  slip  it  into  the  test  tube ;  do  the  same  with  the  manganese  dioxide. 
Shake  the  test  tube  until  the  chemicals  are  thoroughly  mixed ;  then  hold 


Oxygen. 


469 


the  test  tube  in  a  horizontal  position  and  roll  or  shake  it  until  the  mix- 
ture is  spread  along  the  tube  its  entire  length.  Insert  the  stopper  with 
its  tubes,  and  clamp  the  test  tube  to  the  iron  stand,  as  shown  in  the 


FIG.  102.  —  Apparatus  arranged  for  preparing  oxygen. 

figure,  taking  care  not  to  crush  the  tube  ;  the  test  tube  should  incline 
toward  the  trough,  to  prevent  any  water  from  flowing  back  upon  the 
hot  glass. 

Fill  the  pneumatic  trough  with  water  until  the  shelf  is  just  covered. 
Fill  the  bottles/)///  of  water,  cover  each  with  a  piece  of  filter  paper,  in- 
vert them  in  the  trough,  and  remove  the  filter  paper ;  leave  two  bottles 
on  the  shelf  and  three  on  the  bottom.  The  end  of  the  delivery  tube 
should  rest  on  the  bottom  of  the  trough,  just  under  the  hole  in  the  shelf. 

Heat  the  whole  test  tube  gently  with  a  flame  about  8  centimeters  (or  3 
inches)  high.  When  the  gas  bubbles  regularly  through  the  water,  slip  a 
bottle  over  the  hole.  The  gas  will  rise  in  the  bottle  and  force  out  the 
water.  Move  the  flame  slowly  along  the  test  tube,  but  concentrate  the 
heat  toward  the  closed  end,  and  always  keep  the  flame  behind  any  water 
which  may  be  driven  out  of  the  mixture.  If  the  gas  is  evolved  too  rapidly, 
lessen  the  heat;  if  too  slowly,  increase  it ;  if  not  at  all,  examine  the 
stopper  and  the  rubber  connecting  tube  for  leaks,  and  adjust  accordingly. 
When  the  first  bottle  of  gas  is  full,  remove  and  cover  it  with  a  piece  of 
wet  filter  paper,  and  slip  another  bottle  over  the  hole.  When  five  bot- 
tles of  gas  have  been  collected,  remove  the  end  of  the  delivery  tube 
from  the  water,  lest  the  cold  water  be  drawn  up  into  the  hot  test  tube 
as  the  gas  contracts. 

Perform  the  next  experiment  at  once. 


470  Experiments. 

Experiment  6.  — Properties  of  Oxygen. 

Proceed  as  follows  with  the  oxygen  prepared  in  the  preceding 
experiment. 

(a)  Dip  a  glowin^_stickjDf^wood  into  one  bottle,  and  observe  the 
change.  Remove  the  stick,  and  repeat  as  many  times  as  possible. 
Does  the  gas  burn?  How  does  the  glowing  stick  change?  What 
property  of  oxygen  does  this  experiment  show? 

(b}  Put  a  small  piece  of  sulphur  in  the  deflagrating  spoon,  hold 
the  spoon  in  the  flame  until  the  blue  flame  of  the  burning  sulphur  can  be 
seen,  then  lower  the  spoon  into  a  bottle  of  oxygen.  Notice  the  change 
in  the  flame.  Describe  it.  Brush  a  little  of  the  vapor  cautiously 
toward  the  nose.  Of  what  does  the  odor  remind  you?  (Plunge  the 
spoon  into  water  to  extinguish  the  burning  sulphur,  and  covef  the 
bottle  with*  a  piece  of  filter  paper.) 

(c)  Hold  the  charcoal  in  the  flame  long  enough  to  produce  a  faint 
glow,  then  lower  IFmto  a  bottle  of  oxygen.  Describe  the  result. 

(a)  Melt  the  sulphur  in  the  deflagrating  spoon,  and  dip  the  unwound 
end  of  the  wire  picture  cord  into  the  melted  sulphur.  Lower  the  end 
coated  with  burning  sulphur  into  a  bottle  of  oxygen.  The  iron  wire 
should  burn  brilliantly.  Describe  the  change.  Sometimes  the  sub- 
stance produced  by  the  change  coats  the  inside  of  the  bottle  Describe 
it,  if  it  is  visible. 

(tf)   With  the  remaining  bottle,  repeat  any  of  the  above  experiments. 

EXERCISES  : 

(1)  Write  a  brief  account  of  the  above  experiments  in  your  note 
book,  answering  all  questions  and  directions. 

(2)  Sketch  the  apparatus  used  to  prepare  oxygen. 

(3)  Summarize  the  properties  of  oxygen. 

(4)  What  is  its  most  characteristic  property? 

(NOTE.  —  The  test  tube  used  in  Experiment  5  may  be  cleaned  with 
warm  water.) 

Experiment  7.  —  Preparation  of  Oxygen  from  Mercuric  Oxide. 

Materials :  Mercuric  oxide,  stick  of  wood. 

Put  a  little  mercuric  oxide  on  the  end  of  a  narrow  piece  of  paper 
creased  lengthwise,  and  slip  the  powder  into  a  test  tube.  The  pow- 
der should  nearly  fill  the  round  end  of  the  test  tube.  Hold  the  test 
tube  in  a  horizontal  position,  shake  it  to  spread  the  powder  into  a  thin 


Hydrogen. 


471 


layer,  attach  the  test-tube  holder,  and  heat  the  test  tube  (still  horizontal) 
in  the  upper  part  of  the  Bunsen  flame.  Do  not  heat  one  place,  but 
move  the  tube  back  and  forth.  As  soon  as  a  definite  change  is  noticed 
inside  the  tube,  insert  a  glowing  stick  of  wood.  Observe  and  describe 
the  change.  If  there  is  no  change,  heat  strongly,  and  test  again. 
.What  gas  is  liberated?  Observe  the  deposit  inside  the  tube.  What  is 
it?  If  its  nature  is  doubtful,  let  the  tube  cool,  and  examine  again. 

EXERCISES  : 

(1)  Describe  briefly  the  whole  experiment. 

(2)  What  historical  interest  has  this  experiment? 

(NOTE.  —  If  the  test  tube  has  been  partially  melted,  save  it  for  a  sub- 
sequent experiment.) 

HYDROGEN. 

Experiment  8.  —  Preparation  of  Hydrogen.  Materials :  Granu- 
lated zinc,  dilute  sulphuric  acid,  pneumatic  trough,  four  bottles,  filter 
paper,  taper,  matches.  The  apparatus  is  shown  in  Figure  103.  A  is  a 
large  test  tube  provided  with  a  two-hole  stopper,  through  which  passes 
the  safety  tube,  B,  and  the  right-angle  bend,  C\  the  long  (15  cm.  or  6 
in.)  delivery  tube,  E,  is  attached  to  the  bent  tube  by  the  rubber  tube,  D. 
Precaution.  Keep  all  flames  away  from  the  hydrogen  generator. 
Fill  the  test  tube  half  full  of  granulated  zinc  as  follows : 
Crease  a  piece  of  paper  lengthwise,  pour  the  zinc  from  the 
bottle  upon  the  paper,  incline  the  test  tube,  and  slip  the  zinc 
.into  it  from  the  paper  —  do  not  drop  it  in.  Insert  the  stopper 
with  its  tubes ;  if  the  end  of  the  safety  tube  does  not  go  in 
easily,  hold  the  test  tube  in  a  horizontal  position  and  shake 
the  zinc  about,  and  at  the  same  time  push 
the  stopper  gently  but  firmly  into  place. 
Clamp  the  apparatus  into  the  position  shown 
in  the  figure  or  stand  it -in  a  test-tube  rack. 
Fill  the  pneumatic  trough 
with  water  as  before,  and  ad- 
just the  apparatus  so  that  the 
end  of  the  delivery  tube  rests 

on  the  bottom  of  the  trough 
FIG.  103.- Apparatus  for  preparing  hydrogen.      under  ^  ho]e  ^  ^  ^^ 

Fill  the  bottles  with  water  and  invert  them   in  the  trough,  as  in 
Experiment  5. 


O 


472  Experiments. 

Pour  enough  dilute  sulphuric  acid  through  the  safety  tube  to  fill  the 
test  tube  about  half  full,  taking  care  to  leave  a  little  acid  in  the  lower 
bend  of  the  safety  tube.  This  precaution  prevents  the  gas  from  escap- 
ing from  the  back  of  the  apparatus ;  if  at  any  time  the  gas  should  flow 
backward,  pour  a  little  acid  into  the  bend ;  if  the  acid  does  not  flow 
down  the  safety  tube,  loosen  the  stopper  for  an  instant.  As  soon  as 
the  interaction  of  the  zinc  and  sulphuric  acid  produces  hydrogen,  the 
gas  will  bubble  freely  through  the  water  in  the  trough.  Slip  a  bottle 
over  the  hole,  and  collect  and  remove  the  bottle  of  gas  as  in  Experi- 
ment 5,  taking  care  to  cover  the  bottle  firmly  with  a  piece  of  wet  filter 
paper.  If  the  evolution  of  gas  slackens  or  ceases,  add  a  little  more  acid 
through  the  safety  tube.  Collect  four  bottles  of  hydrogen,  and  proceed 
at  once  with  the  next  experiment. 

Experiment  9.  —  Properties  of  Hydrogen. 

Study  as  follows  the  hydrogen  gas  prepared  above :  — 

(a}  Uncover  a  bottle  for  an  instant  to  let  a  little  air  in,  and  then 
drop  a  lighted  match  into  the  bottle.  Describe  the  result. 

(£)  Remove  the  paper  from  a  bottle  of  hydrogen,  and  allow  it  to 
remain  uncovered  for  three  minutes  —  by  the  clock.  Then  show  the 
presence  or  absence  of  hydrogen  by  dropping  a  lighted  match  into  the 
bottle.  Describe  the  result.  What  property  of  hydrogen  is  shown  by 
this  experiment? 

(c}  Verify  your  answer  to  the  last  question,  thus :  Hold  a  bottle  of 
air  over  a  covered  bottle  of  hydrogen,  remove  the  paper,  and  bring  the 
mouths  of  the  bottles  close  together.  (See  Fig.  i.)  Hold  them  there 
for  a  minute  or  two,  then  stand  the  bottles  on  the  desk  and  cover  them 
with  wet  filter  paper.  Drop  a  lighted  match  into  each  bottle.  What 
has  become  of  the  hydrogen?  How  does  (c)  verify  (£)? 

(d}  Invert  a  covered  bottle  of  hydrogen,  remove  the  paper,  and 
quickly  thrust  a  lighted  taper  up  into  the  bottle.  Withdraw  the  taper 
and  then  insert  it  again.  Does  the  hydrogen  burn?  If  so,  where? 
Does  the  taper  burn  when  in  the  bottle?  When  out  of  the  bottle  ? 
Feel  of  the  neck  of  the  bottle ;  describe  and  explain.  What  three 
properties  of  hydrogen  are  shown  by  this  experiment  ? 

Experiment  10.  — Burning  Hydrogen.  (Teacher's  Experi- 
ment.) Materials:  Apparatus  shown  in  Figure  2,  which  consists  of  a 


Hydrogen.  473 

500  cubic  centimeter  flask  fitted  with  a  two-hole  rubber  stopper,  safety 
tube,  and  double  right-angle  bend ;  the  last  is  attached  to  a  U-tube, 
which  is  also  connected  to  a  delivery  tube  provided  with  a  short  piece 
of  capillary  glass  tubing;  calcium  chloride,  small  bottle,  platinum  wire, 
cotton,  granulated  zinc,  dilute  sulphuric  acid. 

Fill  the  U-tube  two  thirds  full  of  calcium  chloride,  put  a  wad  of 
cotton  beneath  the  stopper  of  each  arm,  and  connect  the  U-tube  with 
the  generator  and  the  delivery  tube. 

Stand  the  apparatus  on  the  table,  examine  all  joints  to  be  sure  they 
are  tight,  extinguish  all  flames  in  the  vicinity,  and  proceed  exactly 
according  to  the  following  directions :  — 

Pour  slowly  but  continuously  through  the  safety  tube  enough  (about 
50  cubic  centimeters)  dilute  sulphuric  acid  upon  at  least  25  grams  of 
granulated  zinc  to  produce  a  steady  current  of  hydrogen  gas  for  about 
five  minutes.  It  is  advisable  to  use  considerable  zinc  and  a  moderate 
amount  of  acid.  Acid  must  not  be  added  after  the  evolution  of  gas 
begins,  unless,  of  course,  the  experiment  is  begun  anew.  Let  the  gas 
bubble  through  the  acid  for  at  least  two  minutes  by  actual  observation, 
then  attach  the  capillary  tube  by  the  rubber  connector  to  the  end  of 
the  delivery  tube,  leaving  a  short  space  between  the  ends  of  the  two 
glass  tubes  so  that  the  rubber  tube  may  be  compressed  suddenly,  if 
necessary.  Let  the  gas  run  for  another  full  minute.  This  latter  pre- 
caution is  to  drive  all  air  out  of  the  capillary  tube.  Light  the  hydrogen, 
and  observe  at  once  the  nature  of  the  flame,  its  color,  heat  (by  holding 
a  match  or  platinum  wire  over  it),  and  any  other  striking  property. 
Then  hold  a  small  dry  bottle  over  the  flame  in  such  a  position  that  the 
flame  is  just  inside  the  bottle.  When  conclusive  evidence  of  the  prod- 
uct of  burning  hydrogen  is  seen  inside  the  bottle,  remove  the  bottle, 
and  extinguish  the  flame  at  once  by  pinching  the  rubber  connector. 
Remove  the  generator  to  the  hood,  and  if  the  evolution  of  hydrogen  is 
still  brisk,  dilute  the  acid  by  pouring  water  through  the  safety  tube. 
Examine  the  inside  of  the  bottle.  What  is  the  deposit  ?  Explain  its 
formation. 

EXERCISES  FOR  THE  CLASS: 

(1)  What  does  this  experiment  suggest  about  the  composition  of 
water  ? 

(2)  Does  this  experiment  illustrate  oxidation?     Why?     Synthesis? 
Why? 

(3)  Describe  the  whole  experiment,  and  sketch  the  apparatus. 


474  Experiments. 

WATER. 

Experiment  11.  —  General  Distribution  of  Water.  Materials: 
Wood,  meat,  potato. 

Heat  successively  in  dry  test  tubes  a  small  piece  of  wood,  of  meat, 
or  of  potato  (or  any  other  fresh  vegetable) .  Hold  the  open  end  of  the 
test  tube  lower  than  the  other  end.  Is  there  conclusive  evidence  of 
water?  Since  most  animal  and  vegetable  substances  act  similarly,  what 
general  conclusion  can  be  drawn  from  this  experiment  ? 

Experiment  12.  —  Simple    Tests    for    Impurities    in   "Water. 

Materials:  Distilled  water,  water  containing  dirt,  a  sulphate,  a  chloride, 
and  a  lime  compound ;  nitric  acid,  ammonium  hydroxide,  acetic  acid, 
sulphuric  acid  (concentrated),  solutions  of  potassium  permanganate, 
silver  nitrate,  barium  chloride,  ammonium  oxalate  ;  and  limewater. 

(a)  Organic  Matter.    Fill  a  clean  test  tube  half  full  of  distilled  water, 
and  another  with  water  containing  a  little  dirt  or  a  bit  of  paper.     Add 
to  each  test  tube  a  drop  or  two  of  concentrated  sulphuric  acid  and  suffi- 
cient potassium  permanganate  solution  (made  from  distilled  water)  to 
color  each  liquid  a  light  purple,  as  nearly  alike  as  possible.     Label  one 
tube,  and  then  heat  gently  nearly  to  the  boiling  point  the  tube  contain- 
ing the  impure  water.     As  soon  as  a  definite  change  is  seen,  heat  the 
other  cautiously,  as  too  sudden  heat  may  cause  the  liquid  to  "bump  out." 
Organic  matter  decolorizes  potassium  permanganate  solution.     Which 
tube  shows  the  more  organic  matter? 

(b)  Chlorides.    To  a  test  tube  half  full  of  distilled  water  add  a  few 
drops  of  nitric  acid,  and  then  a  few  drops  of  silver  nitrate  solution.    Do 
the  same  with  water  known  to  contain  a  chloride  in  solution.     What  is 
the  difference  between  the  results  ?     The  cloudiness,  or  solid,  is  due  to 
silver  chloride,  which  is  always  formed  when  silver  nitrate  is  added  to 
hydrochloric  acid  or  a  chloride  in  solution  (chlorides  being  closely  related 
to  hydrochloric  acid).    Silver  chloride  is  soluble  in  ammonium  hydroxide. 
Try  it.     This  is  the  usual  test  for  chlorides  (and  conversely  for  soluble 
silver  compounds),  and  will  hereafter  be  used  without  further  description. 

(c}  Sulphates.  To  a  test  tube  half  full  of  distilled  water  add  a  few 
drops  of  sulphuric  acid  and  a  few  drops  of  barium  chloride  solution. 
The  white  precipitate  is  barium  sulphate.  It  is  insoluble  in  all  common 
liquids,  and  is  always  formed  when  barium  chloride  is  added  to  sulphu- 
ric acid  or  a  sulphate  in  solution  (sulphates  being  closely  related  to  sul- 
phuric acid).  Test  the  impure  water  for  sulphates. 


Water.  475 

(y  )  Lime  Compounds.  Add  a  few  drops  of  a  fresh  solution  of  ammo- 
nium oxalate  to  a  test  tube  half  full  of  clear  limewater.  Limewater  is 
a  solution  of  calcium  hydroxide,  and  the  white  precipitate  formed  is 
calcium  oxalate,  which  is  soluble  in  hydrochloric  acid  but  not  in  acetic 
acid.  Try  it.  This  is  the  test  for  calcium  compounds,  often  called 
"lime"  compounds,  because  lime,  which  is  calcium  oxide,  is  so  well 
known.  Apply  this  test  to  distilled  water  and  to  water  known  to  con- 
tain calcium  compounds,  and  compare  the  two  results. 

(e)    Summarize  briefly  the  whole  experiment. 

(NOTE.  —  If  time  permits,  this  experiment  should  be  applied  by  the 
class  to  water  whose  impurities  are  unknown.) 

Experiment  13.  —  Distillation.  (Teacher's  Experiment.)  Ma- 
terials: Condenser,  etc.,  shown  in  Figure  6,  potassium  permanganate, 
impure  water,  and  solutions  used  in  Experiment  12. 

Fill  the  flask,  C,  half  full  of  water  known  to  contain  the  impurities 
mentioned  in  Experiment  12,  add  a  few  crystals  (3  or  4)  of  potassium 
permanganate,  and  connect  with  the  condenser  as  shown  in  Figure  6. 
Attach  the  inlet  tube  to  the  faucet,  fill  the  condenser  slowly,  and  regu- 
late the  current  so  that  a  small  stream  flows  continuously  from  the 
outlet  tube  into  the  sink  or  waste  pipe.  Heat  the  liquid  in  C  gradually, 
and  when  it  boils,  regulate  the  heat  so  that  the  boiling  is  not  too  vio- 
lent. As  the  distillate  collects  in  the  receiver,  Z?,  test  separate  portions 
for  organic  matter,  chlorides,  sulphates,  and  calcium  compounds. 

EXERCISES  FOR  THE  CLASS  : 

(1)  Is  organic  matter  found  ? 

(2)  Is  mineral  matter  found  ? 

(3)  If  the  distilling  liquid  had  contained  a  volatile  substance,  like 
ammonia  or  alcohol,  would  the  distillate  contain  such  a  substance  ? 

Experiment  14.  —  Solubility  of  Gases. 

(a}  Warm  a  little  faucet  water  in  a  test  tube.  Is  there  immediate 
evidence  of  a  previously  dissolved  gas  ?  Is  there  evidence  of  much 
gas  ?  What  effect  has  increased  heat  ? 

(6)  Warm  slightly  a  few  cubic  centimeters  of  ammonium  hydroxide 
in  a  test  tube.  Do  the  results  resemble  those  in  (a)  ?  As  soon  as  the 
final  result  is  obtained,  pour  the  remaining  liquid  down  the  sink  and 
flush  well  with  water. 


476 


Experiments. 


(V)    Repeat  (£),  using  a  little  concentrated  hydrochloric  acid.     Do 
the  results  resemble  those  of  (a)  and  (b)  ? 
ANSWER  : 

(1)  How  does  increased  temperature  affect  the  solubility  of  gases  ? 

(2)  What  gases  dissolve  freely  in  water  ? 

Experiment  15.  —  Solubility  of  Liquids.  Materials:  Alcohol, 
kerosene,  glycerine,  carbon  disulphide. 

(a)  To  a  test  tube  half  full  of  water  add  a  little  alcohol  and  shake. 
Is  there  evidence  of  solution  ?  Add  a  little  more  and  shake.  Add  a 
third  portion.  Is  there  still  evidence  of  solution  ?  Draw  a  conclusion 
as  to  the  solubility  of  alcohol  in  water. 

(£)  Repeat  («),  using  successively  kerosene,  glycerine,  and  carbon 
disulphide.  Observe  the  results  and  conclude  accordingly. 

(c)    Summarize  the  results  in  a  table. 

Experiment  16.  —  Solubility  of  Solids.  Materials:  About  20 
grams  of  powdered  copper  sulphate,  6  grams  of  powdered  potassium 
chlorate,  i  gram  of  calcium  sulphate. 

(a)  Label  three  test  tubes  I,  II,  III.  Fill  each  about  one  third  full. 
To  I  add  i  gram  of  powdered  copper  sulphate,  to  II  add  i  gram  of 
powdered  potassium  chlorate,  to  III  add  i  gram  of  calcium  sulphate. 
Shake  each  test  tube,  and  then  allow  them  to  stand  undisturbed  for  a 
few  minutes.  Is  there  evidence  of  solubility  in  each  case?  Is  there 
evidence  of  a  varying  degree  of  solubility?  If  III  is  doubtful,  carefully 
transfer  a  portion  of  the  clear  liquid  to  an  evaporating  dish  by  pouring 
it  down  a  glass  rod  (see  Introduction,  §  6  (i )(#)),  and  evaporate  to  dry- 
ness.  Is  there  now  conclusive  evidence  of  solution?  Draw  a  general 
conclusion  from  this  experiment.  Save  solutions  I  and  II  for  (£). 

Tabulate  the  results  of  (d)  as  follows,  using  the  customary  terms  to 
express  the  degree  of  solubility  :  — 

TABLE  OF  SOLUBILITY  OF  TYPICAL  SOLIDS. 


SOLUTE. 

SOLVENT. 

RESULTS. 

i  .   Copper  sulphate 
2.   Potassium  chlorate 

Water  at  tempera- 
ture of  labora- 

I. 

2. 

3.   Calcium  sulphate 

tory. 

3- 

Water.  477 

(£)  Heat  I  and  add  gradually  4  more  grams  of  powdered  copper 
sulphate.  Does  it  all  dissolve?  Heat  II  and  add  4  more  grams  of 
powdered  potassium  chlorate.  Does  it  all,  or  most  all,  dissolve?  What 
general  effect  has  increased  heat  on  the  solubility  of  solids?  What  is 
the  difference  between  this  general  result  and  that  in  Experiment  14? 
Save  the  solutions  for  (c) . 

(c)  Heat  I  and  II  nearly  to  boiling,  and  as  the  temperature  in- 
creases add  the  respective  solids.  Do  not  boil  the  liquid  away.  Is 
there  a  limit  to  their  solubility?  Draw  a  general  conclusion  from  these 
typical  results. 

Experiment  17.  —  Supersaturation.  Material:  Sodium  thio- 
sulphate. 

Fill  a  test  tube  nearly  full  of  crystallized  sodium  thiosulphate  and 
add  a  very  little  water.  Warm  slowly.  As  solution  occurs,  heat 
gradually  to  boiling.  Add  sodium  thiosulphate  until  no  more  will 
dissolve.  Pour  the  solution  into  a  warm,  clean,  dry  test  tube  and 
let  it  stand  until  cool.  Then  drop  in  a  small  crystal  of  sodium  thio- 
sulphate and  watch  for  any  simple  but  definite  change.  What  hap- 
pens? Is  the  excess  of  solid  large?  How  does  a  supersaturated 
solution  differ  from  a  saturated  one? 

Experiment  18.  —  Water  of  Crystallization.  Materials:  Crys- 
tallized sodium  carbonate,  gypsum,  copper  sulphate,  evaporating  dish, 
gauze-covered  ring  (or  tripod). 

(a)  Heat  a  few  small  crystals  of  sodium  carbonate  in  a  dry  test  tube, 
inclining  the  test  tube  so  that  the  open  end  is  the  lower.     What  is  the 
evidence  that  they  contained  water  of  crystallization?     If  there  is  any 
marked  change  in  the  appearance  of  the  crystals,  describe  and  explain  it. 

(b)  Repeat,  using  a  crystal  of  gypsum.     Answer  the  question  asked 
in  (a). 

(c}  Heat  two  or  three  small  crystals  of  copper  sulphate  in  an  evapo- 
rating dish  which  stands  on  a  gauze-covered  ring.  As  the  action  pro- 
ceeds, hold  a  dry  funnel  or  glass  plate  over  the  dish.  Is  there  conclusive 
evidence  of  escaping  water  of  crystallization  ?  Do  the  crystals  change 
in  color?  In  shape?  Can  the  form  of  the  crystals  be  changed  by 
gently  touching  the  mass  with  a  glass  rod?  Continue  to  heat  until  the 
resulting  mass  is  a  bluish  gray.  Let  the  dish  cool.  Meanwhile  heat  a 
test  tube  one  half  full  of  water.  When  the  dish  has  cooled  somewhat, 


478  Experiments. 

pour  the  hot  water  slowly  into  the  dish  upon  the  copper  sulphate.  Ex- 
plain the  change  in  color,  if  any.  If  there  are  any  lumps,  crush  them 
with  a  glass  rod.  Let  the  clear  solution  evaporate  for  several  hours. 
Are  crystals  deposited?  If  not,  heat' a  few  minutes,  and  cool  again. 
If  so,  why  ?  Have  they  water  of  crystallization,  and,  if  so,  where  did 
they  get  it? 

Experiment  19.  —  Efflorescence. 

Put  a  fresh  crystal  of  sodium  carbonate  and  of  sodium  sulphate  on 
a  piece  of  filter  paper,  and  leave  them  exposed  to  the  air  for  an  hour  or 
more.  Describe  any  marked  change.  What  does  this  change  show 
about  the  air  ?  About  the  crystal  ? 

Experiment  20.  —  Deliquescence. 

Put  on  a  glass  plate  or  block  of  wood  a  small  piece  of  granulated 
calcium  chloride  and  of  sodium  hydroxide.  Leave  them  exposed  to 
the  air  for  an  hour  or  more.  Describe  any  marked  change  which  takes 
place.  Compare  the  action  with  that  of  Experiment  19. 

Experiment  21.  —  Solution  and  Chemical  Action.  Materials: 
Powdered  tartaric  acid,  sodium  bicarbonate,  lead  nitrate,  potassium  di- 
chromate,  mortar,  dish  of  water. 

(#)  Mix  in  a  dry  mortar  small  but  equal  amounts  of  powdered  tar- 
taric acid  and  sodium  bicarbonate.  Is  there  any  decided  evidence  of 
chemical  action  ?  Pour  the  mixture  into  a  dish  of  water.  Is  there  con- 
clusive evidence  of  chemical  action  ? 

(b)  Repeat,  using  powdered  lead  nitrate  and  powdered  potassium 
dichromate. 

Describe  the  results  in  (a)  and  (b).  How  does  solution  influence 
chemical  action  ?  Why  are  so  many  solutions  used  in  the  laboratory  ? 

Experiment  22. — Electrolysis  of  Water.  (Teacher's  Experi- 
ment.) Materials :  Hofmann  apparatus,  sulphuric  acid,  taper,  matches, 
short  piece  of  capillary  glass  tubing. 

Fill  the  Hofmann  apparatus,  Figure  10,  with  water  containing  10  per 
cent  of  sulphuric  acid,  so  that  the  water  in  the  reservoir  tube  stands  a 
short  distance  above  the  gas  tubes  after  the  stopcock  in  each  has  been 
closed.  Connect  the  platinum  terminal  wires  with  a  battery  of  at  least 
two  cells.  As  the  action  proceeds,  small  bubbles  of  gas  rise  and  collect 


Water.  479 

at  the  top  of  each  tube.  Allow  the  current  to  operate  until  the  smaller 
volume  of  gas  is  from  8  to  10  centimeters  in  height.  Measure  the 
height  of  each  gas  column.  Assuming  that  the  tubes  have  the  same 
diameter,  the  volumes  are  in  approximately  the  same  ratio  as  their  heights. 
How  do  the  volumes  compare  ? 

Test  the  gases  as  follows :  (#)  Hold  a  glowing  taper  over  the  tube 
containing  the  smaller  quantity  of  gas,  cautiously  open  the  stopcock  to 
allow  the  water  (or  air)  to  run  out  of  the  glass  tip,  and  then  let  out  a 
little  gas  upon  the  glowing  taper.  What  is  the  gas  ?  Repeat  until  the 
gas  is  exhausted.  Care  must  be  taken  not  to  lose  the  gas.  It  is  ad- 
visable to  have  at  hand  several  partially  burned  tapers  or  thin  splints, 
in  case  any  escaping  water  extinguishes  the  first  one.  (t>)  Open  the 
other  stopcock  long  enough  to  force  out  the  water  in  the  glass  tip ; 
close  the  stopcock,  and,  by  means  of  a  short  rubber  tube,  attach  the 
capillary  tube  close  to  the  end  of  the  glass  tip.  Open  the  stopcock 
again,  let  out  the  gas  slowly,  and  hold  at  the  same  time  a  lighted  match 
at  the  end  of  the  tip,  then  immediately  thrust  a  taper  into  the  small 
and  almost  colorless  flame.  What  is  the  gas  ?  Repeat  until  the  gas  is 
exhausted. 

EXERCISES  FOR  THE  CLASS: 

(1)  Describe  the  whole  experiment. 

(2)  Draw  a  general  conclusion  from  this  experiment. 

(3)  What  does  this   experiment   show  about  the  composition  of 
water  ? 

(4)  Sketch  the  apparatus. 

Experiment  23.  —  Interaction  of  Water  and  Chlorine.  (Teach- 
er's Experiment.)  Materials:  Glass\ube  I  meter  long  and  about  2 
centimeters  in  diameter,  cork  for  one  end,  evaporating  dish,  chlorine 
water. 

Construct  a  chlorine  generator,  as  described  in  Experiment  38,  and 
prepare  about  250  cubic  centimeters  of  chlorine  water  by  causing  the 
gas  to  bubble  through  a  bottle  of  water  until  the  water  smells  strongly 
of  the  gas.  Close  one  end  of  the  tube  with  a  cork.  The  cork  must 
fit  air  tight,  and  as  a  precaution  should  be  smeared  (after  insertion) 
with  vaseline  or  coated  with  paraffin.  Fill  the  tube  full  of  chlorine 
water,  cover  the  open  end  with  the  thumb  or  finger,  invert  the  tube,  and 
immerse  the  open  end  in  the  evaporating  dish,  which  should  be  nearly 


480  Experiments. 

full  of  chlorine  water.  Clamp  the  tube  in  an  upright  position,  and  stand 
the  whole  apparatus  where  it  will  receive  the  direct  sunlight  for  at  least 
six  hours.  Bubbles  of  gas  will  soon  appear,  rise,  and  collect  at  the 
top.  When  sufficient  gas  for  a  test  has  collected,  unclamp  the  tube, 
cover  the  open  end  with  the  thumb  or  finger,  invert  the  tube,  and  put  a 
glowing  taper  into  the  gas.  Repeat  as  long  as  any  of  the  gas  remains. 

EXERCISES  FOR  THE  CLASS: 

(1)  What  gas  is  produced  by  the  interaction  of  chlorine  and  water  ? 

(2)  Describe  this  experiment. 

(3)  What  does  it  show  about  the  composition  of  water  ? 

(4)  Sketch  the  apparatus. 

Experiment  24. — Interaction  of  Water  and  Sodium.  Mate- 
rials :  Sodium,  pneumatic  trough  filled  with  water  as  usual,  tea  lead,  for- 
ceps, red  litmus  paper. 

Precaution.  Sodium,  shotdd  be  handled  cautiously  and  used  strictly 
according  to  directions.  Small  fragments  must  not  be  left  about  nor 
thrown  into  the  refuse  jar,  but  into  a  large  vessel  of  water  especially  pro- 
vided for  that  purpose. 

(a)  If  the  sodium  is  brown,  scrape  off  the  coating.     Cut  off  a  piece  of 
sodium  not  larger  than  a  small  pea,  and  drop  it  upon  the  water  in  the 
trough.     Stand  far  enough  away  so  that  you  can  just  see  the  action. 
Wait  until  you  are  sure  the  action  has  stopped,  and  then  describe  all  you 
have  seen. 

(b)  The  action  in  (a)  may  be  further  studied  as  follows  :  Fill  a  test 
tube  with  water,  invert  it,  and  clamp  it  in  the  trough  so  that  the  mouth  is 
over  the  hole  in  the  shelf  of  the  trough.     Wrap  a  small  piece  of  sodium 
loosely  in  a  piece  of  tea  lead  about  5  centimeters  (2  inches)  square,  make 
two  or  three  small  holes  in  the  tea  lead,  and  then  thrust  it  under  the 
shelf  of  the  trough  with  the  forceps.      A  gas  will  rise  into  the  test  tube. 
Proceed  similarly  with  additional  small  pieces  of  sodium  and  dry  tea 
lead  until  the  test  tube  is  nearly  full  of  gas  ;  then  unclamp  and  remove, 
still  keeping  the  tube  inverted.     Hold  a  lighted  match,  for  an  instant,  at 
the  mouth  of  the  tube.     Observe  the  result,  watching  especially  the 
mouth  of  the  tube.     What   is  the  gas?     Why?      Remembering  that 
sodium  is  an  element,  where  must  the  gas  have  come  from?     If  there  is 
any  doubt  about  the  nature  of  the  gas,  collect  more,  and  subject  it  to 
those  tests  which  will  prove  its  nature. 


The  Air.  481 

(V)  Put  a  piece  of  filter  paper  on  the  water  in  the  trough,  and  before 
it  sinks  drop  a  small  piece  of  sodium  upon  it.  Stand  back  and  observe 
the  result.  Wait  for  the  slight  explosion  which  usually  occurs  soon 
after  the  action  stops.  Describe  all  you  have  seen.  What  burned? 
What  caused  it  to  burn?  To  what  is  the  vivid  color  probably  due? 
(In  answering  these  questions,  utilize  your  knowledge  (i)  of  the  prop- 
erties of  the  gases  previously  studied,  and  (2)  of  the  usual  accompani- 
ment of  chemical  action,  suggested  here  by  the  melting  of  the  sodium.) 

(d}  Test  the  water  in  the  trough  with  red  litmus  paper.  Push  the 
paper  to  the  bottom  or  to  the  place  where  it  is  certain  that  chemical 
action  between  water  and  sodium  has  taken  place.  Test  until  the  red 
litmus  paper  has  undergone  a  decided  change  in  color.  Describe  this 
final  result.  With  another  piece  of  red  litmus  paper  test  a  solution  of 
sodium  hydroxide.  Is  the  result  similar?  Dip  a  glass  rod  or  the  plati- 
num test  wire  (see  Int.  §  5  (4))  into  this  solution  and  hold  it  in  the 
Bunsen  flame.  Describe  the  result.  Is  the  color  of  this  flame  and  that 
noticed  in  (<:)  the  same?  Are  the  dissolved  substances  identical? 

(e)  WThat  does  the  whole  experiment  show  about  the  composition  of 
water  ? 

THE  AIR. 

Experiment  25. —  Composition  of  the  Air.  Materials:  Solu- 
tions of  pyrogallic  acid  and  potassium  hydroxide,1  pneumatic  trough  half 
filled  with  water  at  the  temperature  of  the  room,  500  and  25  cubic  centi- 
meter graduated  cylinders.  The  apparatus  consists  of  an  Erlenmeyer 
flask  (250  cubic  centimeters)  provided  with  a  one-hole  rubber  stopper 
into  (but  not  through)  which  passes  a  short  glass  tube ;  to  the  outer  end 
of  this  tube,  which  projects  2.5  centimeters  (i  inch)  above  the  stopper, 
a  rubber  tube  (5  centimeters  or  2  inches  long)  is  tightly  fastened ;  a 
Hofmann  screw  is  attached  to  the  rubber  tube  close  to  the  end  of  the 
glass  tube. 

(a)  The  volume  of  the  flask  is  found  thus  :  Fill  the  flask  completely 
with  water  from  the  pneumatic  trough.  Loosen  the  screw  and  push 
the  stopper  into  the  flask  as  far  as  it  will  go.  Wipe  the  flask  dry  and 
carefully  remove  the  stopper.  Pour  most  of  the  water  from  the  flask 
into  the  500  cubic  centimeter  graduate,  and  read  the  volume :  the  last 
portions  of  the  water  in  the  flask  should  be  poured  into  the  25  cubic 

1  The  pyrogallic  acid  is  a  10  per  cent  solution,  and  the  potassium  hydroxide 
50  per  cent. 


482  Experiments. 

centimeter  graduate,  so  that  the  volume  can  be  read  accurately.     (See 
Fig.  101).     Record  the  total  volume  of  the  flask  as  shown  in  {d}. 

(b)  Measure  exactly  10  cubic  centimeters  of  pyrogallic  acid  in  the 
small  graduate  (see  Int.  §  6  (3)  (d)),  and  pour  it  into  the  flask.     Add 
20  cubic  centimeters  of  potassium  hydroxide  solution,  and  insert  the 
rubber  stopper  quickly  and  firmly.     Tighten  the  screw.     Shake  the 
flask  vigorously  for  a  minute.     Then  invert  it  and  watch  the  surface  of 
the  liquid  for  bubbles.     If  any  appear,   the  apparatus  leaks.      Find 
the  leak,  if  any,  start  the  experiment  again  from  (£),  taking  care   to 
remedy  the  defect  before  the  flask  is  shaken.     If  no  bubbles  appear, 
continue  to  shake  at  intervals  from  fifteen  to  twenty  minutes.     During 
this  operation  the  oxygen  is  absorbed  by  the  solution. 

(c)  Place  the  flask  on  its  side  in  the  water  of  the  pneumatic  trough, 
and  open  the  screw,  taking  care  (i)  not  to  let  any  of  the  solution  run 
out,  (2)  nor  to  let  too  much  water  run  in,  and  (3)  to  keep  the  end  of 
the  rubber  tube  constantly  below  the   surface.     After  the  water  has 
stopped  running  in,  remove  the  flask  from  the  trough.     Open  the  flask, 
put  a  glowing  stick  into  the  gas,  and  observe  the  result.     The  gas  is 
nitrogen.     Measure  carefully  the  volume  of  the  final  liquid  in  the  flask. 

(d)  Record  and  calculate  as  follows  :  — 

(a)  Volume  of  original  solution  =  30  cc. 

(b)  Capacity  of  flask  =       cc. 

(c)  Volume  of  air  taken  (b  —  a)  =  , 

(d)  Final  volume  of  liquid  = 

(e)  Volume  of  water  which  entered  (d  —  a) 
(f  )  Per  cent  of  water  which  entered  (e  -s-  c) 

But  the  per  cent  of  entering  water  equals  the  per  cent  of  gas  ab- 
sorbed, hence 

(g)  Per  cent  of  oxygen 

(h)  Per  cent  of  nitrogen  (100  —  g)  = 

Experiment  26.  —  Air  contains  Water  Vapor. 

Prove  by  an  experiment  that  air  contains  water  vapor. 

Experiment  27.  —  Air  contains  Carbon  Dioxide. 

(a)  Expose  a  small  bottle  of  limewater  to  the  air.  After  a  short 
time,  examine  the  surface  of  the  liquid.  Describe  the  change.  Ex- 
plain it. 


Acids,  Bases,  and  Salts.  483 

(£)  If  a  blast  lamp  (or  bicycle  pump)  is  available,  replace  the  lamp 
with  a  glass  tube,  and  force  air  through  a  bottle  half  full  of  limewater, 
until  a  definite  change  occurs.  Describe  it.  Explain  it. 

ACIDS,  BASES,  AND  SALTS. 

Experiment  28.  —  General  Properties  of  Acids.  Materials : 
Dilute  sulphuric,  nitric,  and  hydrochloric  acids,  glass  rod,  litmus  paper 
(both  colors),  zinc. 

Fill  separate  test  tubes  one  third  full  of  each  of  the  acids.  Label  the 
tubes  in  some  distinguishing  manner. 

(a)  Dip  a  clean  glass  rod  into  each  acid  and  cautiously  taste  it. 
Describe  the  taste  by  a  single  word. 

(b)  Dip  a  clean  glass  rod  into  each  acid  and  put  a  drop  on  both  kinds 
of  litmus  paper.    The  striking  change  is  characteristic  of  acids  ;  draw  a 
general  conclusion  from  it. 

(£)  Slip  a  small  piece  of  zinc  into  each  test  tube  successively.  If  no 
chemical  action  results,  warm  gently.  Test  the  most  obvious  product 
by  holding  a  lighted  match  inside  of  each  tube.  What  gas  comes  from 
the  hydrochloric  and  sulphuric  acids? 

(d}  Summarize  the  general  results  of  this  experiment. 

Experiment  29.  —  General  Properties  of  Bases.  Materials: 
Litmus  paper  (both  colors),  glass  rod,  sodium  hydroxide  and  potassium 
hydroxide  solutions,  and  ammonium  hydroxide. 

(a)  Rub  a  little  of  each  liquid  between  the  fingers,  and  describe  the 
feeling.    Cautiously  taste  each  liquid  by  touching  to  the  tip  of  the  tongue 
a  rod  moistened  in  each,  and  describe  the  result. 

(b)  Test  each  solution  with  litmus  paper.     Describe  the  result. 

(c)  Summarize  the  general  results  of  this  experiment. 

(d)  Compare  acids  and  bases  as  to  taste  and  to  reaction  with  litmus. 

Experiment  30.  —  A  Property  of  Many  Salts  and  All  Neutral 
Substances.  Materials:  Litmus  paper  (both  colors),  glass  rod,  dilute 
solutions  of  sodium  chloride,  potassium  nitrate,  potassium  sulphate,  and 
barium  chloride. 

Test  each  solution  with  litmus  paper.     Describe  the  result.     Com- 
pare with  the  litmus  reaction  of  acids  and  bases. 
Draw  a  general  conclusion  from  this  experiment. 


484 


Experiments. 


Experiment  31.  —  The  Nature  of  Common  Substances. 

Determine  by  the  litmus  test  the  nature  of  lemon  juice,  vinegar, 
sweet  and  sour  milk,  washing  soda,  borax,  wood  ashes,  faucet  water, 
baking  soda,  sugar,  cream  of  tartar,  the  juice  of  any  ripe  fruit  and  any 
green  fruit. 

Make  a  solution  of  each  of  the  solids  before  testing.  Tabulate  the 
results  as  follows  :  — 

NATURE  OF  COMMON  SUBSTANCES. 


ACID. 


ALKALINE. 


NEUTRAL. 


Experiment  32.  —  Neutralization.  Materials:  Sodium  hydrox- 
''ide  (solid),  hydrochloric  acid,  nitric  acid,  silver  nitrate  solution,  blue 
litmus  paper,  glass  rod,  evaporating  dish,  gauze-covered  ring. 

Dissolve  a  small  piece  of  sodium  hydroxide  in  an  evaporating  dish 
half  full  of  water.  Slowly  add  dilute  hydrochloric  acid,  until  a  drop 
taken  from  the  dish  upon  a  glass  rod  reddens  blue  litmus  paper.  Then 
evaporate  to  dryness  by  heating  over  a  piece  of  wire  gauze  supported 
by  a  ring.  Since  the  residue  mechanically  holds  traces  of  the  excess 
of  hydrochloric  acid  added,  it  is  necessary  to  remove  this  acid  before 
applying  any  test.  Heat  the  dish  until  all  the  yellow  color  disappears, 
then  moisten  the  residue  carefully  with  a  few  drops  of  warm  water  and 
heat  again  to  remove  the  last  traces  of  acid.  This  precaution  is  essen- 
tial to  the  success  of  the  experiment. 

Test  a  portion  of  the  residue  with  litmus  paper  to  find  whether  it  has 
acid,  alkaline,  or  neutral  properties.  Taste  a  little.  Test  (a}  a  solu- 
tion of  the  residue  for  a  chloride,  and  (b}  a  portion  of  the  solid  residue 
for  sodium.  (See  Exps.  12  (t>)  and  24.)  Draw  a  definite  conclusion 
from  the  total  evidence. 


Heat,  Light,  Electricity,  and  Chemical  Action.     485 

HEAT,   LIGHT,   ELECTRICITY,  AND   CHEMICAL  ACTION. 

Experiment  33.  —  Heat  and  Chemical  Action.  Materials : 
Lime,  evaporating  dish,  match. 

Put  a  small  piece  of  lime  in  an  evaporating  dish,  and  sprinkle  a  little 
water  over  it.  Watch  for  a  change.  If  no  marked  change  soon  occurs, 
add  a  little  more  water.  Describe  the  change.  Touch  a  match  to  the 
mass.  Is  there  evidence  of  much  heat?  What  caused  the  heat? 

Experiment  34.  —  Heat  and  Chemical  Action.  Materials:  Sul- 
phur, powdered  iron,  dilute  hydrochloric  acid. 

Put  about  3  grams  of  sulphur  and  3  grams  of  powdered  iron  in  a 
test  tube.  Cover  the  mouth  of  the  test  tube  with  the  thumb  and 
shake  until  the  two  substances  are  well  mixed.  Attach  the  test  tube  to 
the  holder  and  heat  strongly  in  the  flame.  As  soon  as  the  sulphur 
melts  and  boils  and  the  contents  give  evidence  of  decided  chemical 
action,  remove  the  test  tube  at  once  from  the  flame,  and  watch  the 
change.  Is  there  evidence  of  heat?  Of  increasing  heat?  Of  much 
heat? 

When  the  tube  is  cool,  break  the  end,  and  examine  the  contents.  De- 
scribe it.  It  is  a  compound  called  iron  sulphide,  and  is  the  product  of 
the  chemical  action  which  was  started  by  heat.  But  the  chemical  action 
itself  was  so  vigorous  that  it  increased_the  heat. 

The  fact  that  the  product  differs  from  the  original  mixture  may  be 
shown  as  follows  :  Add  dilute  hydrochloric  acid  to  a  part  of  the  product 
and  also  to  a  little  of  the  original  mixture,  testing  the  gaseous  product 
in  each  case  by  the  odor.  Is  the  odor  the  same? 

State  briefly  how  heat  and  chemical  action  are  related,  using  this 
experiment  as  an  illustration. 

Experiment  35.  —  Light  and  Chemical  Action.  Materials  : 
Potassium  bromide,  silver  nitrate  solution,  funnel,  filter  paper,  glass  rod. 

Dissolve  a  crystal  of  potassium  bromide  in  a  test  tube  one  fourth  full 
of  water,  add  an  equal  volume  of  silver  nitrate  solution,  and  shake.  The 
precipitate  is  silver  bromide.  Describe  it.  Filter  (see  Int.  §  4).  Remove 
the  filter  paper  from  the  funnel,  unfold  it,  and  expose  the  silver  bromide 
for  a  few  minutes  to  the  light  —  sunlight,  if  possible.  Describe  the 
change.  What  caused  the  change?  How  is  this  property  of  silver 
bromide  utilized  ? 


486  Experiments. 

Experiment  36.  —  Electricity  and  Chemical  Action.  (Teacher's 
Experiment.) 

Repeat  Experiment  22. 
EXERCISES  FOR  THE  CLASS: 

(1)  Define  electrolysis,  electrode,  electrolyte,  ion,  anion,  cation. 

(2)  State  briefly  the  accepted  explanation  of  the  electrolysis  of  water. 

(3)  Is  hydrogen  an  anion  or  cation?     At  what  electrode  does  it 
collect  ? 

(4)  Answer  the  same  questions  (as  in  3)  about  oxygen. 

Experiment  37.  —  Electricity  and  Chemical  Action.  (Teach- 
er's Experiment.)  Materials  :  Starch,  potassium  iodide,  mortar  and 
pestle,  filter  paper,  sheet  tin  (or  iron),  battery  of  two  or  more  cells. 

Grind  together  in  a  mortar  a  lump  of  starch  and  a  crystal  of  potas- 
sium iodide.  Add  enough  water  to  make  a  thin  liquid.  Dip  a  strip  of 
filter  paper  into  the  mixture,  and  spread  the  wet  paper  upon  a  sheet  of 
tin  (or  iron) .  Press  the  end  of  the  wire  attached  to  the  zinc  (of  the 
battery)  upon  the  tin,  and  draw  the  other  wire  across  the  sheet  of 
paper.  The  marks  are  caused  by  iodine  which  is  liberated  from  the 
potassium  iodide  and  colors  the  starch. 

EXERCISES  FOR  THE  CLASS: 

(1)  Describe  briefly  this  experiment. 

(2)  Iodine  is  a  non-metal.     At  what  electrode  is  it  liberated?     Is 
iodine  an  anion  or  a  cation  ? 

CHLORINE. 

{Do  not  inhale  chlorine?) 

Experiment  38.  —  Preparation  of  Chlorine.  Materials:  Con- 
centrated hydrochloric  acid,  30  grams  manganese  dioxide,  bundle  of 
fine  brass  wire,  strip  of  calico,  paper  with  writing  in  lead  pencil  and  in 
ink,  litmus  paper  (both  colors),  taper.  The  apparatus  is  shown  in  Fig- 
ure 104.  It  is  the  same  as  that  used  to  prepare  hydrogen ;  and  there 
are  also  needed  four  bottles,  a  wooden  block  (about  10  centimeters  or 
4  inches  square)  with  a  hole  in  the  center,  and  four  glass  plates  to 
cover  the  bottles. 

Weigh  the  manganese  dioxide  upon  a  piece  of  paper  creased  length- 
wise. Slip  it  into  the  test  tube,  A  (see  Int.  §  6  (£)).  Arrange  the  appa- 


Chlorine. 


487 


ratus  as  shown  in  the  figure.     Pour  enough  concentrated  hydrochloric 

acid  through  the  safety  tube  to  cover  the  man- 
ganese  dioxide.      Heat  gently  with   a  small 

flame,  keeping  the  flame  below  the  level  of  the 

contents  of  the  test  tube.     Chlorine  is  rapidly 

evolved  as  a  greenish  gas,  and  passes  into  the 

bottle,  G,  which  should  be  removed  when  full 

(as  seen  by  the  green  color)  and  covered  with 

a  glass  plate  ;  the  bottle  may  be  easily  removed 

by  holding  the  block,  F,  in   one   hand  and 

pulling  the  bottle,  G,  aside,  bending  the  whole 

delivery  tube  at  the  same  time  at  the  rubber 

connection,    D.       If    the    evolution    of   gas 

slackens,  add   more   acid   through   the  safety 

tube.     Collect  four  bottles,  and   perform  the 

next  experiment  at  once.  FIG.  104.  —  Apparatus 

arranged    for     preparing 
chlorine. 
Experiment  39.  —  Properties  of  Chlorine. 

Study  as  follows  the  gas  prepared  above :  — 

(a)  Heat  the  bundle  of  brass  wire  and  thrust  it  into  a  bottle  of  chlo- 
rine.    Describe  the  result,  especially  the  evidence  of  chemical  action 
and  of  new  products. 

(b)  Into  a  bottle  of  dry  chlorine  put  a  piece  of 
calico,  litmus  paper  (both  colors),  and  paper  contain- 
ing writing  in  black  and  in  red  ink.  Allow  the  whole 
to  remain  undisturbed  for  a  few  minutes  and  then 
describe  the  change,  if  any.  Add  several  drops  of 
water,  and  describe  the  change.  Draw  a  general  con- 
clusion from  the  whole  experiment. 

(c}  Hold  a  burning  taper  in  a,  bottle  of  chlorine 
long  enough  to  observe  the  result.    Draw  a  conclusion. 
Verify  it  thus  :    Fold  a  strip  of  filter  paper  (about   10 
centimeters  or  4  inches  wide)  into  the  shape  shown 
FIG.  105.— Fluted   ln  Figure  105;  cautiously  heat1  about  10  cubic  centi- 
paper.  meters   of  turpentine  in  a  large  test  tube;   saturate 


1  Hold  the  test  tube  with  the  holder.      Remember  that  turpentine  ignites  easily. 
If  the  turpentine  catches  fire,  press  a  damp  towel  over  it. 


488  Experiments. 

the  paper  with  the  hot  turpentine  and  drop  it  into  a  bottle  of  chlorine. 
Describe  the  result.  When  the  action  is  over,  examine  the  paper,  and 
draw  a  conclusion  regarding  the  action  between  hot  turpentine  and 
chlorine. 

Wax  (in  the  taper)  and  turpentine  are  mainly  compounds  of  hydro- 
gen and  carbon.  Explain  the  result  in  (c) . 

ANSWER  : 

(1)  Many  metals  act  like  the  brass  in  (#).     What  general  conclu- 
sion can  be  drawn  about  the  reaction  of  chlorine  and  metals? 

(2)  What  is  essential  for  the  bleaching  action  of  chlorine? 

(3)  What  does  (c}  show  about  the  attraction  between  chlorine  and 
hydrogen  ? 

(4)  What  class  of  chemical  changes  is  illustrated  by  (#)  ?      What 
classes  by  (c)  ? 

(5)  What  class  of  chemical  changes  is  illustrated  by  the  preparation 
of  chlorine? 

(6)  What  three  striking  properties  has  chlorine?      How  can  it  be 
distinguished  from  all  gases  previously  studied? 

Experiment  40.  —  Bleaching  by  Bleaching  Powder.  Mate- 
rials :  Bleaching  powder,  sulphuric  acid,  calico. 

Put  a  little  bleaching  powder  into  a  test  tube  and  add  enough  water 
to  make  a  thin  paste.  Add  a  few  drops  of  dilute  sulphuric  acid,  and 
then  dip  a  strip  of  bright-colored  calico  into  the  mixture.  Remove  the 
calico  in  a  few  minutes,  and  wash  it  with  water.  Describe  the  change 
in  the  calico. 

Experiment  41.  —  Preparation  of  Hydrochloric  Acid.  Mate- 
rials :  The  apparatus  used  in  Experiment  38 ;  20  grams  sodium  chlo- 
ride, concentrated  sulphuric  acid,  pneumatic  trough  filled  with  water  as 
usual,  stick  of  wood,  litmus  paper  (blue),  ammonium  hydroxide. 

(a)  Put  8  cubic  centimeters  of  water  in  a  small  bottle  or  evaporating 
dish,  cautiously  add  12  cubic  centimeters  of  concentrated  sulphuric  acid, 
and  stir  until  the  two  are  mixed.  While  this  mixture  is  cooling,  weigh 
the  salt,  slip  it  into  the  test  tube,  and  then  arrange  the  apparatus  as 
shown  in  Figure  104.  Pour  half  the  cold  acid  mixture  through  the 
safety  tube,  let  it  settle  through  the  salt,  and  then  add  the  remaining 
acid.  Heat  gently  with  a  low  flame,  as  in  the  preparation  of  chlorine. 
Hydrochloric  acid  gas '  is  evolved,  and  passes  into  the  bottle,  which 


Chlorine.  489 

should  be  removed  when  full,  as  directed  under  chlorine.  A  piece  of 
moist  blue  litmus  paper  held  at  the  mouth  of  the  bottle  will  show  when 
it  is  full.  Collect  these  bottles,  cover  each  with  a  glass  plate,  and  set 
aside  until  needed. 

(b)  As  soon  as  the  third  bottle  of  gas  has  been  collected,  removed, 
and  covered,  put  in  its  place  a  bottle  one  fourth  full  of  water.     Adjust 
its  height  (if  necessary)  by  wooden  blocks   so   that  the   end  of  the 
delivery  tube  is  just  above  the  surface  of  the  water.     Continue  to  heat 
the  generator  at  intervals,  and  the  gas  will  be  absorbed  by  the  water. 
Shake  the  bottle  occasionally. 

Meanwhile  study  the  gas  already  collected. 

Experiment  42.  — Properties  of  Hydrochloric  Acid  Gas. 

Proceed  as  follows  with  the  hydrochloric  acid  gas  prepared  by 
Experiment  41  :  — 

(#)  Insert  a  blazing  stick  of  wood  into  a  bottle.  Remove  as  soon  as 
the  change  is  noticed.  Describe  the  change.  Compare  the  action  with 
the  behavior  of  hydrogen  and  of  oxygen  under  similar  conditions. 

(^)  Hold  a  piece  of  wet  filter  paper  near  the  mouth  of  the  same 
bottle.  Describe  the  result.  What  is  the  cause? 

(c)  Invert  a  bottle,  and  stand  it  upon  the  shelf  of  the  pneumatic 
trough.     Describe  any  change  noticed  inside  the  bottle  after  a  few 
minutes.     What  property  of  the  gas  does  the  result  illustrate  ?     Verify 
the   observation   by  a  simple    test    applied   to   the   contents   of    the 
bottle. 

(d)  Drop  into  the  remaining  bottle  of  gas  a  piece  of  filter  paper  wet 
with  ammonium   hydroxide.     Describe   the  result.     What   name  has 
the  product? 

(e)  State  other  properties  of  hydrochloric  acid  gas  which  you  have 
observed  ;  e.g.  color,  odor,  density. 

Proceed  at  once  with  the  next  experiment. 

Experiment  43.  —  Properties  of  Hydrochloric  Acid. 

Remove  the  bottle  in  which  the  hydrochloric  acid  gas  is  being  ab- 
sorbed (see  Exp.  41  (£)),  and  study  the  solution  as  follows :  — 

(a)  Determine  its  general  properties,  e.g.  taste  (cautiously),  action 
with  litmus,  and  with  zinc. 

(^)  Add  to  a  test  tube  half  full  of  the  hydrochloric  acid  a  few  drops 
of  nitric  acid  and  of  silver  nitrate  solution.  The  white,  curdy  precipitate 


49° 


Experiments. 


is  silver  chloride.  Filter  part  of  the  contents  of  the  test  tube,  and  ex- 
pose the  precipitate  to  the  sunlight.  Describe  the  change  which  soon 
occurs.  To  the  remaining  contents  of  the  test  tube  add  ammonium 
hydroxide,  and  shake.  Describe  the  result. 

Experiment  44.  —  Tests  for  Hydrochloric  Acid  or  a  Chloride. 

(a)  What  is  a  simple  test  for  hydrochloric  acid  gas  or  for  concen- 
trated hydrochloric  acid  ? 

{b}    What  is  the  usual  test  for  hydrochloric  acid  ? 

(c)  Dissolve  a  little  sodium  chloride  in  a  test  tube  half  full  of  water, 
and  apply  the  test  designated  in  (£).  (Suggestions.  See  Exps.  12  (b) 
and  43  (£).) 

COMPOUNDS  OF  NITROGEN. 

Experiment  45.  —  Preparation  of  Ammonia.  Materials  :  1 5  grams 
lime,  15  grams  ammonium  chloride,  3  bottles,  2  glass  plates,  pneu- 
matic trough  filled  as  usual,  litmus  paper,  stick  of 
wood,  filter  paper.  The  apparatus  is  shown  (in 
part)  in  Figure  106.  The  large  test  tube,  A,  is 
provided  with  a  one-hole  rubber  stopper  to  which 
is  fitted  the  right-angle  bend,  C,  connected  with  a 
short  glass  tube,  B  (12  centimeters  or  5  inches 
long),  by  the  rubber  tube,  D. 

(a)  Weigh  the  lime  and  ammonium  chloride 
separately,  mix  them  thoroughly  on  a  piece  of 
paper,  and  slip  the  mixture  into  the  test  tube  to 
which  a  little  water  has  been  previously  added. 
Add  a  little  water.  Quickly  insert  the  stopper 
with  its  tubes,  and  clamp  the  test  tube  as  shown 
in  the  figure  (taking  care  not  to  crush  the  test 
tube) . 

Slip  the  glass  delivery  tube,  B,  into  a  bottle, 
invert  the  bottle,  and  hold  it  so  that  the  tube  is 
in  the  position  shown  in  the  figure.  Heat  the 
test  tube  gently  with  a  low  flame,  beginning  near  the  top  of  the  mix- 
ture and  gradually  working  downward.  Ammonia  gas  will  pass  up 
into  the  bottle,  which  should  be  removed  when  full  and  covered  with 
a  glass  plate.  A  piece  of  moist  red  litmus  paper  held  near  the  mouth 
will  show  when  the  bottle  is  full.  Do  not  smell  at  the  mouth  of  the 
bottle.  Collect  two  bottles  and  set  aside  until  needed. 


FIG.  106.  —  Appara- 
tus for  preparing  and 
collecting  ammonia  gas. 


Compounds  of  Nitrogen.  491 

(b)  As  soon  as  the  last  bottle  has  been  collected,  rearrange  the  appa- 
ratus to  absorb  the  ammonia  gas  in  water,  as  in  the  case  of  hydrochloric 
acid  (see  Exp.  41  (^)).     Replace  the  short  glass  tube  by  the  delivery 
tube,  E,  which  should  pass  through  the  wooden  block,  fy  into  a  bottle, 
G,  one  fourth  full  of  water,  so  that  the  end  is  just  above  the  surface  of  the 
water  (see  Fig.  104).     Continue  to  heat  the  generator  at  intervals,  and 
the  gas  will  be  absorbed  by  the  water.     Shake  the  bottle  occasionally. 

While  the  solution  is  being  prepared,  study  the  gas  already  collected. 

Experiment  46. —  Properties  of  Ammonia  Gas. 

Proceed  as  follows  with  the  ammonia  gas  prepared  in  Experiment 
45  :- 

(a)  Test  the  gas  in  one  bottle  with  moist  litmus  paper  and  with  a 
blazing  stick.  Describe  the  result.  Compare  the  action  with  the 
behavior  of  hydrogen,  oxygen,  and  hydrochloric  acid  gas,  under  similar 
circumstances. 

(£)  Invert  the  same  bottle  and  stand  it  upon  the  shelf  of  the  pneu- 
matic trough.  Describe  any  change  noticed  inside  the  bottle.  What 
property  of  the  gas  is  revealed  ?  Is  it  a  marked  property  ?  Test  the 
contents  of  the  bottle  with  litmus  paper  (both  colors). 

(c)  Pour  a  few  drops  of  concentrated   hydrochloric  acid  into  an 
empty,  warm,  dry  bottle.    Roll  the  bottle  until  the  inside  is  well  coated. 
Cover  it  with  a  glass  plate,  invert  it,  and  stand  it  upon  a  covered  bottle 
of  ammonia  gas.     Remove  both  plates  at  once,  and  hold  the  bottles 
together  by  grasping  them  firmly  about  their  necks.      Describe  the 
action,  giving  all  the  evidence  of  chemical  action.      What  is  the  white 
product  ? 

Experiment  47. —  Properties  of  Ammonium  Hydroxide. 

Remove  the  bottle  in  which  the  ammonia  gas  is  being  absorbed 
(see  Exp.  45,  (£)),  and  study  the  resulting  ammonium  hydroxide  as 
follows : —  , 

(a)  Determine    the   general   properties,  e.g.  taste  and   odor  (cau- 
tiously), feeling,  action  with  litmus. 

(b)  Warm  a  little  in  a  test  tube.     What  gas  is  evolved? 

(c)  Try  the  effect  of  ammonium  hydroxide  on  a  grease  spot.   Describe 
the  result. 

Experiment  48. —  Neutralization  of  Ammonia.  Materials: 
Ammonium  hydroxide,  hydrochloric  acid,  evaporating  dish,  sodium 
hydroxide  solution,  litmus  paper,  gauze-covered  ring. 


492  Experiments. 

Fill  an  evaporating  dish  one  fourth  full  of  ammonium  hydroxide,  and 
slowly  add  dilute  hydrochloric  acid,  stirring  constantly,  until  the 
solution  is  just  neutral  or  faintly  acid.  Evaporate  to  dry  ness,  very 
slowly,  on  a  gauze-covered  ring.  Test  the  residue  as  follows  : — 

(#)    Is  it  an  acid,  alkali,  or  salt  ? 

(^)  Warm  a  little  with  sodium  hydroxide  solution.  What  is  formed? 
Draw  a  conclusion  as  to  the  nature  of  the  residue. 

(c}  Support  the  dish  on  the  gauze  and  warm  gently  until  a  decided 
change  occurs.  Describe  the  result.  What  compound  do  the  fumes 
suggest? 

(</)  Verify  the  observations  and  conclusions  by  repeating  (3)  and 
(c)  with  ammonium  chloride  from  the  laboratory  bottle. 

(Y)  What  is  the  main  product  of  the  interaction  of  ammonium 
hydroxide  and  hydrochloric  acid? 

Experiment  49. —  Preparation  of  Nitric  Acid.  Materials: 
Glass  stoppered  retort,  sand  bath  pan  and  sand,  bottle,  30  grams  sodium 
nitrate,  concentrated  sulphuric  acid,  funnel. 

Weigh  the  sodium  nitrate  and  slip  it  into  the  retort  (  see  Int.  §  6  (i) 
(c}  ).  Attach  the  retort  by  a  clamp  to  an  iron  stand  so  that  (i)  Us  bulb 
rests  on  the  sand  bath,  supported  by  a  ring,  and  (2)  the  end  of  its  neck 
passes  into  an  inclined  bottle  which  rests  on  the  table.  The  nitric  acid 
which  is  generated  in  the  bulb  will  pass  down  the  neck  and  condense, 
partly  in  the  neck  and  partly  in  the  bottle.  The  bottle  should  be  partially 
covered  with  a  piece  of  wet  filter  paper,  especially  where  the  neck  of  the 
retort  enters.  It  is  advisable,  though  not  always  necessary,  to  place  a 
block  of  wood  against  the  bottom  of  the  bottle  to  keep  it  in  the  desired 
position. 

Slip  a  funnel  through  the  tubulure  of  the  retort  as  far  as  it  will  reach, 
and  pour  the  acid  through  the  funnel  into  the  retort.  Remove  the  funnel 
and  insert  the  stopper  of  the  retort  tightly.  Heat  gently.  Brown 
fumes  will  appear  in  the  retort,  and  nitric  acid  will  pass  into  the  receiver. 
Distil  at  as  low  a  temperature  as  possible,  as  long  as  any  nitric  acid  runs 
down  the  neck  of  the  retort. 

Pour  the  nitric  acid  into  a  test  tube  or  small  bottle  for  use  in  Experi- 
ment 50. 

Allow  the  contents  of  the  retort  to  cool,  add  a  little  warm  water,  let 
the  whole  stand  until  the  contents  are  loosened,  and  then,  pour  into  a 
bottle  for  use  in  Experiment  53. 


Compounds  of  Nitrogen.  493 

Experiment  50.  —  Properties  of  Nitric  Acid.  Materials :  Quill 
toothpick,  indigo  solution.  Add  twice  its  volume  of  water  to  the  nitric 
acid  prepared  in  Experiment  49,  and  proceed  as  follows :  — 

(a)  Boil  a  piece  of  a  quill  toothpick  in  a  portion  of  this  diluted  nitric 
acid.     How  is  the  quill  changed  at  first?     What  is  the  effect  of  contin- 
ued heating?     Pour  off  the  acid,  and  wash  the  quill  with  water.     Is  the 
color  permanent  ? 

(b)  Add  a  dozen  or  more  drops  of  nitric  acid  to  a  dilute  solution  of 
indigo.     Describe   the   change.     Will   ammonium    hydroxide   restore 
the  original  color?     Is  the  change  temporary  or  permanent?     What,  in 
all  probability,  is  the  general  character  of  the  change  —  combination  or 
decomposition?     Draw  a  general  conclusion  from  (a)  and  (£)  regard- 
ing the  action  of  nitric  acid  on  organic  matter,  which  is  typified  by  the 
quill  and  indigo. 

EXERCISES  : 

(1)  What  color  has  nitric  acid? 

(2)  Examine  a  bottle  of  nitric  acid  which  has  been  standing  in  the 
laboratory.     What  can  be  said  of  the  stability  of  nitric  acid? 

(3)  State  other  properties  of  nitric  acid  you  have  observed. 

Experiment  51.  —  Test  for  Nitric  Acid  and  Nitrates.  Materials  : 
Concentrated  nitric  and  sulphuric  acids,  ferrous  sulphate,  sodium  nitrate. 

To  a  test  tube  one  fourth  full  of  water  add  a  little  concentrated  nitric 
acid  and  shake.  Add  an  equal  volume  of  concentrated  sulphuric  acid. 
Shake  until  the  acids  are  well  mixed,  then  cool  by  holding  the  test  tube 
in  running  water.  Make  a  cold,  dilute  solution  of  fresh  ferrous  sulphate 
and  pour  this  solution  carefully  down  the  side  of  test  tube  upon  the 
nitric  acid  mixture.  Where  the  two  solutions  meet,  a  brown  or  black 
layer  will  appear,  consisting  of  a  compound  formed  by  the  interaction 
of  the  nitric  acid  and  the  ferrous  sulphate.  It  is  an  unstable  com- 
pound and  will  often  decompose  if  the  test  tube  Is  shaken.  Record 
the  observation. 

This  test  is  also  used  for  a  nitrate.  Try  it  with  a  solution  of  sodium 
nitrate.  Record  the  result. 

Experiment  52.  —  A  Special  Test  for  Nitrates.  Materials :  Char- 
coal, block  of  wood,  potassium  nitrate. 

Heat  a  piece  of  charcoal  in  the  Bunsen  flame,  lay  it  on  a  block  of 
wood  or  an  iron  pan,  and  cautiously  sprinkle  powdered  potassium  ni- 


494  Experiments. 

trate  upon  the  hot  surface.  Stand  back  when  the  action  begins.  Ob- 
serve and  describe  the  action,  especially  its  violence  and  rapidity,  also 
the  color  of  the  flame,  the  effect  on  the  charcoal,  and  any  other  charac- 
teristic result.  This  kind  of  chemical  action  is  called  deflagration. 
What  causes  it? 

Experiment  53.  —  The  Solid  Product  of  the  Interaction  of  So- 
dium Nitrate  and  Sulphuric  Acid.  Materials :  Residue  from  Ex- 
periment 49,  evaporating  dish,  glass  rod,  gauze-covered  ring,  distilled 
water,  barium  chloride  solution,  ferrous  sulphate,  concentrated  sulphuric 
acid. 

Pour  the  solid  residue  obtained  in  Experiment  49. into  an  evaporating 
dish,  and  evaporate  to  dry  ness  over  a  piece  of  wire  gauze  in  the  hood. 
As  the  mass  approaches  pasty  consistency,  lessen  the  heat  to  avoid 
spattering.  When  the  mass  is  dry,  heat  strongly  as  long  as  white, 
choking  fumes  are  evolved.  This  last  operation  is  done  to  remove  all 
traces  of  sulphuric  acid,  and  to  complete  the  chemical  change.  Allow 
the  dish  to  cool  gradually,  and  when  cool,  dissolve  some  of  the  white 
solid  in  distilled  water  and  test  separate  portions  for  a  sulphate  and  ni- 
trate (see  Exps.  12  (V)  and  51).  Which  is  it?  Test  another  portion 
for  sodium  (see  Exp.  24  (</)  ).  What  is  the  name  of  the  white  sub- 
stance ? 

Draw  a  general  conclusion  regarding  the  chemical  action  which  oc- 
curs in  the  preparation  of  nitric  acid  by  the  interaction  of  sulphuric 
acid  and  sodium  nitrate. 

Experiment  54.  —  Interaction  of  Nitric  Acid  and  Metals.  Mate- 
rials :  Zinc,  copper,  tin,  iron,  concentrated  nitric  acid. 

Stand  four  test  tubes  in  the  test-tube  rack,  and  slip  into  each  a  few 
small  pieces  of  one  of  the  following  metals  :  zinc,  copper,  tin,  and  iron. 
Add  to  each  test  tube  in  succession  enough  concentrated  nitric  acid  to 
cover  the  metal.  Observe  the  changes,  particularly  (i)  the  vigor  of  the 
action,  (2)  the  nature  and  properties  of  the  products,  especially  color 
and  solubility,  and  (3)  evidence  of  the  evolution  of  hydrogen.  Tabu- 
late these  observations. 

Experiment  55.  —  Interaction  of  Nitric  Acid  and  Copper,  and 
Study  of  Nitric  Oxide  and  Nitrogen  Peroxide. — Materials:  10 
grams  copper  (borings  or  fine  pieces  of  sheet  metal),  concentrated 
nitric  acid,  pneumatic  trough  filled  as  usual,  three  bottles,  three  glass 


Compounds  of  Nitrogen.  495 

plates,  matches,  piece  of  wire  ( 1 5  centimeters  or  6  inches  long)  ;  and 
the  apparatus  used  in  Experiment  8. 

Arrange  the  apparatus  as  in  Experiment  8,  after  putting  the  copper 
into  the  test  tube  (see  Fig.  103).  Insert  the  stopper  tightly,  adjust 
the  delivery  tube,  fill  three  bottles  with  water,  and  invert  them  in  the 
trough.  Pour  just  enough  concentrated  nitric  acid  through  the  safety 
tube  into  the  flask  to  cover  the  copper,  taking  care  to  seal  the  bend  of 
the  safety  tube  with  acid.  Dense  brown  fumes  are  evolved.  If  the 
action  is  too  vigorous,  add  a  little  water  through  the  safety  tube.  Col- 
lect three  bottles  of  the  gas  which  bubbles  from  the  delivery  tube. 
Cover  them  with  glass  plates  and  stand  them  aside  until  needed.  Pour 
the  blue  liquid  in  the  test  tube  into  an  evaporating  dish,  and  evapo- 
rate slowly  to  crystallization  (not  to  dryness)  on  a  gauze-covered  ring  in 
the  hood.  The  crystals,  after  being  dried  between  filter  paper,  should 
be  preserved  in  a  well-stoppered  bottle. 

While  the  solution  is  evaporating,  study  the  gas  as  follows  :  — 

(a)   Observe  its  general  properties  while  covered. 

(b}  Uncover  a  bottle.  Describe  the  result.  Is  the  brown  gas  iden- 
tical in  color  with  the  one  observed  in  the  generator  at  the  beginning 
of  the  experiment? 

(c)  Uncover  a  bottle,  pour  in  about  25  cubic  centimeters  of  water, 
cover  with  the  hand  and  shake  vigorously,  still  keeping  the  bottle 
covered.  Why  has  the  brown  gas  disappeared?  Uncover  the  bottle 
for  an  instant,  then  cover  and  shake  again.  Is  the  result  the  same? 
Repeat,  if  the  result  is  not  definite,  or  does  not  agree  with  previous 
observations. 

(d}  With  the  third  bottle  determine  whether  the  two  gases  will 
burn  or  support  combustion.  A  convenient  flame  is  a  burning  match 
fastened  to  a  stiff  wire.  Plunge  it  quickly  to  the  bottom  at  first  and 

gradually  raise  it  into  the  brown  gas. 

• 
ANSWER  : 

(1)  What  is  the  source  of  the  colorless  gas?    What  is  its  name? 
What  is  the  name  of  the  brown  gas? 

(2)  What  is  the  general  chemical  relation  of  the  two  gases  to  each 
other?     To  the  air  ? 

(3)  Why  is  not  the  brown  gas  collected  in  the  bottles  by  displace- 
ment of  water? 

(4)  Will  either  gas  burn  or  support  combustion? 


496  Experiments. 

(5)  Which  gas  has  been  observed  before?     In  what  experiment? 

(6)  What  is  the  general  relation  of  these  gases  to  nitric  acid? 
Study  the  properties  of  the  crystals  by  determining :  — 

(a)    Solubility  in  water  (cold  and  hot). 
(£)    Action  of  heat. 

(c}    Action  of  their  solution  upon  an  iron  nail. 
(d)    Action  of  their  solution  when  added  to  ammonium  hydroxide. 
(tf)    Presence  of  a  nitrate. 

Compare  the  observed  properties  with  those  of  copper  nitrate  ob- 
tained from  the  laboratory  bottle.  Are  the  two  substances  identical? 

Experiment  56.  —  Preparation  and  Properties  of  Nitrous  Oxide. 

Materials :  Ammonium  nitrate,  pneumatic  trough  filled,  as  usual,  with 
warm  water,  three  bottles,  three  glass  plates,  sulphur,  deflagrating 
spoo'n,  stick  of  wood.  The  apparatus  is  shown  in  Figure  107.  The 
parts  lettered  A,  C,  D,  E  have  been  used  before ;  B+  F,  G,  H  are 
exactly  the  same  as  A,  C,  D,  E  respectively.  (See  page  504.) 

Construct  and  arrange  the  apparatus  as  shown  in  the  figure.  Fill 
the  large  test  tube,  A,  about  half  full  of  ammonium  nitrate.  The  large 
test  tube,  B,  remains  empty.  The  end  of //rests  on  the  bottom  of  the 
pneumatic  trough  as  usual.  It  is  desirable,  though  not  absolutely 
necessary,  to  fill  the  trough  and  bottles  with  warm  water.  Be  sure  the 
apparatus  is  gas-tight. 

Heat  A  gently  with  a  low  flame  (5  centimeters  or  2  inches).  Adjust 
the  apparatus  if  it  leaks.  The  ammonium  nitrate  melts  and  appears  to 
boil.  Regulate  the  heat  so  that  the  evolution  of  the  nitrous  oxide  will  be 
slow.  Notice  the  fumes  which  collect  in  A,  and  the  liquid  which  col- 
lects in  B.  Prepare  three  bottles  of  nitrous  oxide,  free  from  air,  cover- 
ing each  with  a  glass  plate  as  soon  as  removed  from  the  trough.  When 
the  last  bottle  has  been  collected  and  covered,  remove  the  end  of  the 
delivery  tube  from  the  trough. 

Test  the  gas  as  follows  :  — 

(a)  Allow  a  bottle  to  remain  uncovered  for  a  few  seconds.     How 
does  nitrous  oxide  differ  from  nitric  oxide  ? 

(b)  Thrust  a  glowing  stick  of  wood  into  the  same  bottle  of  gas. 
Describe  the  result.     Is  the  gas  combustible?     Does  it  support  com- 
bustion ? 

(c)  The  observations  in  (£)  suggest  that  the  gas  is  oxygen,  but  it  is 
not,  though  this  fact  is  not  easily  proved  by  a  single  experiment.     Put 


Compounds  of  Nitrogen.  497 

a  small  piece  of  sulphur  in  a  deflagrating  spoon,  light  it,  and  lower  the 
burning  sulphur  at  once  into  another  bottle  of  gas.  If  the  experiment 
is  conducted  properly,  the  sulphur  will  not  burn  so  brightly  as  it  would 
in  a  bottle  of  oxygen. 

(cT)  Stand  the  other  bottle  mouth  downward  in  the  pneumatic 
trough,  or  better,  in  a  vessel  of  cold  water.  Describe  the  result.  If 
the  result  is  not  conclusive,  fill  the  bottle  half  full  of  water,  cover  with 
the  hand,  and  shake.  Does  this  observation  help  distinguish  the  gas 
from  oxygen? 

What  in  all  probability  is  the  other  product  (seen  in  B)  of  the  chemi- 
cal change  in  this  experiment?  Could  it  have  been  an  impurity  in  the 
ammonium  nitrate?  What  are  the  fumes  noticed  in  A? 

How  would  you  distinguish  ammonium  nitrate  from  all  other  nitrates  ? 
How  would  you  distinguish  nitrous  oxide  from  (#)  the  other  oxides  of 
nitrogen,  (6)  air,  (c)  oxygen,  {d}  hydrogen,  (e)  nitrogen,  (/)  carbon 
dioxide? 

Experiment  57.  — Preparation  and  Properties  of  Sodium  Nitrite. 

Materials:  10  grams  sodium  nitrate,  20  grams  lead,  iron  sand  bath 
pan,  glass  rod. 

Heat  the  mixture  of  lead  and  sodium  nitrate  on  the  sand  bath  pan, 
which  stands  on  the  ring  of  an  iron  stand.  Stir  the  melted  mass  with 
a  glass  rod.  Some  of  the  lead  will  disappear  and  a  yellowish  brown 
powder  will  be  seen  in  the  molten  mass.  The  action  should  proceed 
until  most  of  the  lead  has  disappeared.  Allow  the  mass  to  cool,  trans- 
fer to  a  mortar,  pulverize,  add  hot  water,  and  filter  the  clearer  portion  ; 
add  more  hot  water  to  the  residue,  and  filter  this  portion.  This  oper- 
ation extracts  the  sodium  nitrite.  Add  to  the  combined  filtrates  several 
drops  of  concentrated  sulphuric  acid.  Describe  the  result.  How  does 
the  result  compare  with  the  action  of  concentrated  sulphuric  acid  on 
sodium  nitrate?  The  yellowish  product  is  lead  oxide.  What  general 
chemical  change  led  to  its  formation?  How  must  the  nitrate  have  been 
changed  ? 

Experiment  58. — Aqua  Regia.  Materials:  Gold  leaf,  concentrated 
nitric  and  hydrochloric  acids,  glass  rod. 

Touch  a  small  piece  of  gold  leaf  with  the  end  of  a  moist  glass  rod, 
and  wash  the  gold  leaf  into  a  test  tube  by  pouring  a  few  cubic  centi- 
meters of  concentrated  hydrochloric  acid  down  the  rod.  Heat  gently 
until  the  acid  just  begins  to  boil.  Does  the  gold  dissolve?  Wash  an- 


498  Experiments. 

other  piece  of  gold  leaf  into  another  test  tube  with  concentrated  nitric 
acid,  and  heat  as  before.  Does  the  gold  dissolve?  Pour  the  contents 
of  one  tube  into  the  other,  and  warm  gently.  Does  the  gold  dissolve? 
Draw  a  conclusion. 

ANSWER  : 

(1)  What  is  the  literal  meaning  and  significance  of  the  term  aqua 
regia  ? 

(2)  What  other  metals  does  aqua  regia  dissolve? 

(3)  What  is  the  chemical  action  of  aqua  regia  on  gold? 

(4)  Upon  what  property  of  nitric  acid  does  the  action  of  aqua  regia 
depend? 

CARBON.      . 

Experiment  59.  —  Distribution  of  Carbon.  Materials:  Hessian 
crucible,  sand,  wood,  cotton,  starch,  sugar,  glass  tube  (or  rod),  candle, 
block  of  wood. 

(a)  Cover  the  bottom  of  a  Hessian  crucible  with  a  thin  layer  of  sand. 
Put  on  the  sand  a  small  piece  of  wood,  a  small,  compact  wad  of  cotton, 
and  a  lump  of  starch.  Fill  the  crucible  loosely  with  dry  sand,  and 
slip  it  into  the  ring  of  an  iron  stand.  Heat  with  a  flame  which  extends 
just  above  the  bottom  of  the  crucible  until  the  smoking  ceases  (approxi- 
mately 20  minutes) .  After  the  crucible  has  cooled  sufficiently  to  handle, 
pour  the  contents  out  upon  a  block  of  wood  or  an  iron  pan.  Examine 
the  contents.  What  is  the  residue?  What  is  hereby  shown  about  the 
distribution  of  carbon  ? 

While  the  crucible  is  heating,  do  the  following :  — 

(ft)  Heat  about  i  gram  of  sugar  in  an  old  test  tube  until  the  vapors 
cease  to  appear.  What  is  the  most  obvious  product? 

(<:)  Close  the  holes  at  the  bottom  of  a  lighted  Bunsen  burner,  and 
hold  a  glass  tube  in  the  upper  part  of  the  flame  long  enough  for  a  thin 
deposit  to  form.  Examine  it,  name  it,  and  state  its  source. 

(d)  Hold  a  glass  tube  in  the  flame  of  a  candle  which  stands  on  a 
block  of  wood,  and  compare  the  result  with  that  in  (<:). 

Draw  a  general  conclusion  regarding  the  distribution  of  carbon. 

Experiment  60.  —  Decolorizing  Action  of  Charcoal.  Materials: 
Animal  charcoal,  indigo  solution,  filter  paper  and  funnel. 

Fill  a  test  tube  one  fourth  full  of  powdered  animal  charcoal  as  follows  : 
Fold  a  narrow  strip  of  smooth  paper  so  that  it  will  slip  easily  into  the 


Carbon. 


499 


test  tube ;  place  the  powder  at  one  end  of  the  troughlike  holder,  slowly 
push  the  paper  into  the  test  tube,  holding  both  tube  and  paper  in  a 
horizontal  position ;  now  hold  the  tube  upright,  and  the  powder  will 
slip  from  the  paper.  Add  10  cubic  centimeters  of  indigo  solution,  shake 
thoroughly  for  a  minute,  and  then  warm  gently.  Filter  through  a  wet 
filter  paper  into  a  clean  test  tube.  Compare  the  color  of  the  filtrate 
with  that  of  the  indigo  solution.  Explain  the  change  in  color. 

Other  organic  substances  besides  indigo  are  similarly  changed. 
Draw  a  general  conclusion  regarding  the  decolorizing  power  of  char- 
coal. 

Experiment  61. — Deodorizing  Action  of  Charcoal.  Materials: 
Wood  charcoal,  hydrogen  sulphide  solution,  test  tube,  and  cork. 

Smell  of  a  weak  solution  of  hydrogen  sulphide  gas.  Fill  a  test  tube 
half  full  of  powdered  wood  charcoal  as  in  Experiment  60,  add  a  little 
hydrogen  sulphide  solution,  and  cork  securely.  If  the  tube  leaks,  make 
the  opening  gas-tight  with  vaseline.  Shake  thoroughly.  After  fifteen 
or  twenty  minutes,  remove  the  stopper  and  smell  of  the  contents.  Is  the 
odor  much  less  offensive?  Repeat,  unless  a  definite  result  is  obtained. 
Explain  the  change. 

Experiment  62.  —  Preparation  of  Carbon  Dioxide.  Materials: 
Lumps  of  marble,  sand,  concentrated  hydrochloric  acid,  stick  of  wood, 
candle  fastened  to  a  wire,  limewater,  four  bottles.  Use  the  same 
apparatus  as  in  the  preparation  of  hydrogen  (see  Exp.  8). 

Cover  the  bottom  of  the  test  tube  with  sand,  add  a  little  water,  and 
carefully  slip  into  it  half  a  dozen  small  lumps  of  marble.  Arrange  the 
apparatus  to  collect  the  gas  over  water,  as  previously  directed.  Add 
through  the  safety  tube  just  enough  concentrated  hydrochloric  acid  to 
cover  the  marble.  Collect  four  bottles,  cover  with  glass  plates  or  wet 
filter  paper,  and  stand  aside  till  needed. 

Allow  the  action  in  the  flask  to  continue,  and  preserve  the  contents 
for  Experiment  64. 

Proceed  at  once  to  the  next  experiment. 

Experiment  63.  —  Properties  of  Carbon  Dioxide. 

Study  the  properties  of  carbon  dioxide  gas  as  follows  :  — 
(a)  Plunge  a  burning  stick  several  times  into  one  bottle.     Describe 
the  result. 


500  Experiments. 

(b}  Lower  a  lighted  candle  into  a  bottle  of  air,  and  invert  a  bottle 
of  carbon  dioxide  over  it,  holding  the  bottles  mouth  to  mouth.  Describe 
the  result.  What  does  this  result  show  about  the  density  of  carbon 
dioxide? 

(V)  Pour  a  little  limewater  into  a  bottle  of  carbon  dioxide,  cover 
with  the  hand,  and  shake  vigorously.  Describe  and  explain  the 
result. 

(d)  Fill  a  bottle  of  carbon  dioxide  one  third  full  of  water,  cover 
tightly  with  the  hand,  and  shake  vigorously.  Invert,  still  covered,  in 
the  pneumatic  trough.  Does  the  result  reveal  any  facts  about  the 
solubility  of  carbon  dioxide  ? 

EXERCISES  : 

(1)  Describe  the  preparation  of  carbon  dioxide. 

(2)  What  do  (a)  and  (£)  show  about  the  relation  of  carbon  dioxide 
to  combustion? 

(3)  What  is  the  test  for  carbon  dioxide? 

(4)  What  chemical  changes  occur  in  the  test  for  carbon  dioxide? 

Experiment  64. —  The  Solid  Product  of  the  Interaction  of 
Calcium  Carbonate  and  Hydrochloric  Acid. 

Filter  the  contents  of  the  test  tube  into  an  evaporating  dish,  adding 
a  little  warm  water  beforehand,  if  the  contents  are  solid.  Evaporate  to 
dryness  in  the  hood  over  a  free  flame  as  long  as  much  liquid  remains. 
As  the  residue  approaches  pasty  consistency,  add  a  little  water,  stand 
the  dish  on  a  gauze-covered  support,  and  move  the  lighted  burner 
underneath.  Heat  the  residue  until  no  fumes  of  hydrochloric  acid  are 
evolved.  Dissolve  some  of  the  residue  in  distilled  water  and  test 
portions  for  (a)  a  chloride  and  (<£)  a  calcium  compound  (see 
Exp.  12  ($),  (</)).  If  a  calcium  compound  is  found,  confirm  the  obser- 
vation thus  :  Dip  a  clean,  moist  platinum  test  wire  (see  Int.  §  5  (4)  ) 
into  the  solid  residue,  and  hold  it  in  the  Bunsen  flame.  If  calcium  is 
present,  the  flame  will  be  colored  a  yellowish  red. 

What  is  the  residue?     Verify  the  conclusion  by  a  simple  experiment. 

Experiment  65.  —  Carbon  Dioxide  and  Combustion.  Materials : 
Limewater,  glass  tube,  candle  attached  to  wire,  stick  of  wood,  two 
bottles. 

(a)  Exhale  through  a  glass  tube  into  a  test  tube  half  full  of  lime- 
water.  Describe  and  explain  the  result. 


Carbon.    .  501 

(£)  Lower  a  lighted  candle  into  a  bottle  and  allow  it  to  burn  for  a  few 
minutes.  Remove  the  candle,  pour  a  little  limewater  into  the  bottle, 
and  shake  vigorously.  Describe  and  explain  the  result. 

(c)  Allow  a  stick  of  wood  to  burn  for  a  short  time  in  a  bottle  (not 
the  one  used  in  (<£)),  and  then  proceed  as  in  (£).  Describe  the  result. 
Does  it  confirm  the  results  obtained  in  (a)  and  (£)  ? 

(d*)  Repeat  (<:),  using  a  piece  of  paper  in  place  of  the  wood.  De- 
scribe the  result.  Does  it  confirm  the  results  obtained  in  (a),  (£),  and 

(0? 

ANSWER  : 

1 i )  What  is  the  source  of  the  carbon  dioxide  in  (a)  ? 

(2)  What  is  one  of  the  gases  escaping  from  chimneys?     From  a 
burning  lamp  ? 

Experiment    66. — Carbonic    Acid.       (Teacher's    Experiment.) 

Materials:  Solutions  of  sodium  hydroxide  and  phenolphthalein,  bottle, 
and  the  carbon  dioxide  generator  used  in  Experiment  62. 

Construct  and  arrange  the  carbon  dioxide  generator  as  in  Experiment 
62.  Fill  the  bottle  nearly  full  of  water,  add  a  few  drops  of  a  solution  of 
phenolphthalein  *  and  just  enough  sodium  hydroxide  solution  to  color 
the  liquid  a  faint  magenta.  Allow  a  slow  current  of  carbon  dioxide 
to  bubble  through  the  liquid  in  the  bottle,  until  a  definite  change  is 
produced  in  the  absorbing  liquid.  Describe  and  explain  it. 

Experiment  67. — Preparation  and   Properties   of    Carbonates. 

Materials:   Marble,   sand,  concentrated  hydrochloric  acid,  limewater. 
The  apparatus  is  the  same  as  that  used  in  Experiment  62. 

(a)  Prepare  a  carbon  dioxide  generator  as  in  Exp.  62,  and  attach  it  by 
a  clamp  to  an  iron  stand  so  that  the  end  of  the  delivery  tube  reaches  to 
the  bottom  of  a  bottle  half  full  of  limewater.     Pass  the  gas  slowly  into 
the  limewater  until  considerable  precipitate  is  formed.      Remove  the 
bottle  and  let  the  precipitate  settle. 

(b)  Meanwhile  pass  carbon  dioxide  slowly  for  about  five  minutes  into 
a  test  tube  nearly  full  of  a  dilute  solution  of  sodium  hydroxide. 

(<:)  Examine  the  precipitate  from  (a)  as  follows :  Pour  off  most  of 
the  liquid  without  disturbing  the  solid  (see  Int.  §  6  (i)  (#)).  Dip  a  glass 
tube  into  limewater,  remove  it,  and  a  drop  will  adhere  to  the  end. 

iThis  compound  is  magenta  in  alkaline  solutions  and  colorless  in  acid  solutions. 


Experiments. 

Pour  a  little  hydrochloric  acid  into  the  test  tube,  shake,  and  hold  the 
"limewater  tube  "  in  the  escaping  gas.  Observe  the  change  in  the  drop 
of  limewater.  If  no  change  occurs,  add  more  acid  to  the  precipitate. 
What  is  the  liberated  gas?  What  is  the  precipitate?  How  was  the 
latter  formed? 

(d)  Proceed  as  in  (c)  with  the  solution  obtained  in  (V).  What  is 
the  liberated  gas?  From  what  compound  did  it  come?  How  was  this 
compound  formed?  How  does  it  differ  from  the  one  formed  in  (a)  ? 

ANSWER  : 

(1)  What  is  the  test  for  a  carbonate? 

(2)  How  may  limewater  be  distinguished  from  a  solution  of  sodium 
or  potassium  hydroxide  ? 

Experiment  68.  —  Detection  of  Carbonates.  Materials :  Hydro- 
chloric acid,  limewater,  glass  tube  ;  baking  soda,  washing  soda,  baking 
powder,  native  chalk,  tooth  powder,  white  lead,  whiting,  old  mortar  (or 
plaster). 

Put  a  little  of  each  of  the  above  solids  in  separate  test  tubes,  add  a 
little  water  and  dilute  hydrochloric  acid,  and  shake  ;  hold  the  "  lime- 
water  tube  "  in  the  escaping  gas,  as  in  Experiment  67.  If  the  action  is 
not  marked,  warm  the  test  tube.  Describe  the  result  in  each  case. 

Experiment  69.  —  Acid  Calcium  Carbonate.  Materials :  Lime- 
water  and  the  carbon  dioxide  generator  used  in  Experiment  62. 

Pass  carbon  dioxide  into  a  test  tube  half  full  of  limewater  until  the 
precipitate  disappears.  Filter,  if  the  liquid  is  not  perfectly  clear,  and 
then  heat.  Describe  the  change.  Explain  the  three  changes  which 
take  place  in  the  test  tube. 

Experiment  70.  —  Preparation  and  Properties  of  Carbon  Mo- 
noxide. (Teacher's  Experiment.)  Materials :  Oxalic  acid,  concen- 
trated sulphuric  acid,  limewater,  pneumatic  trough  filled  as  usual,  three 
bottles,  three  glass  plates.  The  apparatus  is  shown  in  Figure  107 
(p.  504). 

Precaution.  Carbon  monoxide  and  oxalic  acid  are  poisonous.  Hot 
sulphuric  acid  is  dangerous.  Perform  this  experiment  with  unusual 
care. 

Put  10  grams  of  oxalic  acid  in  the  large  test  tube,  A,  and  add  25 
cubic  centimeters  of  concentrated  sulphuric  acid.  Put  enough  lime- 
water  in  B  to  cover  the  end  of  the  tube,  E.  The  end  of  H  should  rest 


Carbon.  503 

on  the  bottom  of  the  pneumatic  trough  just  beneath  the  hole  in  the 
shelf.  Heat  the  tube,  A,  gently,  and  carbon  monoxide  will  be  evolved. 
A  small  flame  must  be  used,  because  the  gas  is  rapidly  evolved  as  the 
heat  increases.  It  is  advisable  to  remove  or  lower  the  flame  as  bubbles 
appear  in  the  tube,  B,  —  regulate  the  heat  by  the  effervescence.  Collect 
all  the  gas,  but  do  not  use  the  first  bottle,  covering  the  bottles  with 
glass  plates  as  they  are  filled,  and  setting  them  aside  temporarily. 
When  the  last  bottle  has  been  collected  and  covered,  loosen  the  stop- 
per in  B,  remove  the  end  of  H  from  the  water  in  the  trough,  and  if  gas 
is  still  being  evolved,  stand  the  whole  apparatus  in  the  hood. 
Test  the  gas  thus  :  — 

(a)  Notice  that  it  is  colorless. 

(b)  Hold  a  lighted  match  at  the  mouth  of  a  bottle  for  an  instant. 
Note  the  flame,  especially  its  color  and  how  it  burns.     After  the  flame 
has  disappeared,  drop  a  lighted  match  into  the  bottle.     Describe  the 
result.     Draw  a  conclusion  and  verify  it  by  (<:). 

(c}  Burn  another  bottle  of  gas,  and  after  the  flame  has  disappeared 
pour  limewater  into  the  bottle  and  shake.  Describe  the  result. 

EXERCISES  FOR  THE  CLASS: 

(1)  What  gas  besides  carbon  monoxide  was  produced,  as  shown 
by£? 

(2)  Summarize  the  observed  properties  of  carbon  monoxide. 

(3)  What  is  the  chemical  relation  of  the  two  oxides  of  carbon? 

(4)  How  can  the  two  oxides  be  changed  into  each  other?    What 
two  general  processes  do  the  changes  illustrate  ? 

Experiment  71. — Preparation  and  Properties  of  Ethylene. 
(Teacher's  Experiment.)  Materials:  Alcohol,  concentrated  sulphuric 
acid,  sand,  pneumatic  trough  filled  as  usual,  two  bottles,  limewater. 
The  apparatus  is  that  used  in  Experiment  70.  f 

Precaution.  A  mixture  of  ethylene  and  air  explodes,  if  ignited. 
Hot  sulphuric  acid  is  dangerous.  Guard  against  flames,  leaks,  and 
breakage. 

Put  5  cubic  centimeters  of  water  in  a  test  tube  and  slowly  pour  upon 
it  15  cubic  centimeters  of  concentrated  sulphuric  acid.  Cool  the  acid 
by  holding  the  test  tube  in  a  stream  of  cold  water.  Put  5  to  7  cubic 
centimeters  of  alcohol  in  the  test  tube,  A,  add  a  little  clean  sand,  and 


5°4 


Experiments. 


then  slowly  pour  in  the  cold  acid.    The  test  tube,  B,  remains  empty.    A 
dish  should  stand  under  A  to  catch  the  contents,  in  case  of  accident. 

Adjust  the  apparatus  as  shown  in  Figure 
107,  taking  care  not  to  crush  the  test  tubes. 
Heat  the  test  tube,  A,  gently  between  the 
bottom  and  the  surface  of  the  contents ,  to 
detect  any  leaks  in  the  apparatus.     Readjust, 
if  necessary.      Heat  gently  to  drive  out  the 
air,  and  when  it  is  judged  that  the  gas  which 
is   being   evolved    is    ethylene,   collect   two 
bottles.     As  the  heat  increases,  the  mixture 
is     apt     to     froth     or 
"bump";     sometimes 
the  gas  is  evolved  sud- 
denly.   Hence  the  heat 
must   be   so  regulated 
that   the  evolution  of 
gas  is  slow.     Especial 


FIG.  107.  —  Apparatus  for  preparing  ethylene. 


care  must  be  taken  not 
to  heat  the  test  tube 
above  the  surface  of 
the  contents,  otherwise 


a  sudden  movement  of  the  hot  liquid  might  crack  the  test  tube.  As  soon 
as  the  gas  has  been  collected,  remove  the  tube,  //,  from  the  water,  and 
if  the  ethylene  is  still  being  evolved,  stand  the  apparatus  in  the  hood. 
When  the  tube,  A,  is  cool  enough  to  handle,  pour  the  contents  down 
the  sink  or  into  a  receptacle  especially  provided  for  dangerous  mixtures. 
Test  the  gas  by  holding  a  lighted  match  at  the  mouth  of  a  bottle. 
Observe  and  record  the  color  and  temperature  of  the  flame,  its  luminos- 
ity, rapidity  of  combustion,  visible  products,  and  any  other  character- 
istic properties.  Repeat  with  the  other  bottle,  and  carefully  observe 
properties  needing  confirmation.  Add  a  little  limewater  to  one  of  the 
bottles  in  which  the  gas  was  burned,  shake,  and  explain  the  result. 
What  evidence  does  this  experiment  present  regarding  the  composition 
of  ethylene? 

Experiment  72. —Preparation  and  Properties    of    Acetylene. 

Fill  a  test  tube  nearly  full  of  water,  stand  the  test  tube  in  a  rack,  and 
drop  two  or  three  very  small  pieces  of  calcium  carbide  into  the  test 


Carbon. 


5°5 


tube.  Acetylene  is  evolved.  After  the  action  has  proceeded  long 
enough  to  expel  the  air,  light  the  gas  by  holding  a  lighted  match  at  the 
mouth  of  the  tube.  Observe  and  record  the  nature  of  the  flame,  espe- 
cially its  color,  intensity,  visible  products  (if  any),  temperature,  etc. 
Hold  a  cold  glass  plate  or  bottle  over  the  flame.  What  does  the  result 
suggest  about  the  composition  of  acetylene?  What  other  evidence  of 
its  composition  is  revealed  by  the  properties  of  the  flame  ? 

Experiment  73.  —  Preparation  and  Properties  of  Illuminating 
(Coal)  Gas.  (Teacher's  Experiment.)  Materials:  Soft  coal,  asbestos, 
pneumatic  trough  filled  as  usual,  three  bottles,  litmus  paper,  filter  paper, 
lead  acetate  (or  nitrate)  solution.  The  apparatus  is  shown  in  Figure  108. 
A  A'  is  an  ignition  tube  from  10  to  15  centimeters  (4  to  6  inches)  long.  A 
spiral  of  copper  wire  is  placed  near  A',  and  the  tube  is  supported  by  a 


FlG.  108.  —  Apparatus  for  preparing  illuminating  gas  from  soft  coal. 

clamp  between  the  wire  and  the  end  of  the  tube.  An  empty  test  tube 
or  bottle  is  connected  with  the  combustion  tube  by  a  bent  tube  passing 
to  the  bottom  of  B  ;  this  vessel  retains  tarry  matter,  which  comes  from 
the  ignition  tube.  The  U-tube  contains  moistened  pink  litmus  paper 
in  the  limb  C,  and  a  narrow  strip  of  filter  paper  moistened  with  a  lead 
compound  (nitrate  or  acetate)  in  the  limb  C',  the  latter  serving  to  detect 
hydrogen  sulphide.  The  bottle,  D,  which  maybe  ar>y  convenient  size, 
is  connected  as  shown  in  the  figure,  and  is  to  be  one  third  full  of  lime- 
water.  The  tube,  ZT,  is  to  be  connected  with  a  delivery  tube  passing  into 
a  pneumatic  trough  arranged  to  collect  a  gas  over  water. 

Fill  AA'  two  thirds  full  of  coarsely  powdered  soft  coal,  which  should 
be  held  in  place  with  a  loose  plug  of  shredded  asbestos.  See  that  all 
connections  are  gas-tight  by  heating  the  ignition  tube  gently ;  if  the 
ipparatus  is  tight,  the  expanded  air  will  bubble  through  the  bottle  D. 
Readjust,  if  necessary. 


506  Experiments. 

Heat  the  whole  ignition  tube  gently  at  first,  and  gradually  increase 
the  heat,  but  avoid  heating  either  end  very  hot,  otherwise  the  closed 
end  may  soften  and  burst  or  the  rubber  stopper  may  melt.  As  the  heat 
increases,  watch  for  marked  changes  in  B,  CC',  and  D.  As  soon  as  the 
slow  bubbling  shows  that  all  air  has  been  driven  out  of  the  apparatus, 
collect,  as  previously  directed,  two  bottles  of  the  gas  evolved.  Cover 
the  bottles  with  wet  filter  paper  as  soon  as  they  are  removed  from  the 
trough.  When  the  last  bottle  has  been  removed,  disconnect  the  ap- 
paratus at  any  convenient  point  between  A'  and  C.  Let  the  ignition 
tube  cool. 

Test  the  gas  by  holding  a  lighted  match  near  the  mouth  of  a  bottle. 
Observe  and  record  the  color  and  heat  of  the  flame.  Is  smoke  formed  ? 
Repeat  with  the  remaining  bottle,  and  observe  more  closely  any  facts 
suggested,  but  not  clearly  shown,  by  the  first  observations. 

Examine  the  contents  of  the  ignition  tube.  Does  it  resemble  coke 
or  some  form  of  carbon  ?  Examine  the  bottle,  B,  for  tarry  matter. 
Does  the  paper  in  C  show  the  formation  of  ammonia  ?  If  the  paper  in 
C  is  black  or  brown,  it  is  caused  by  lead  sulphide,  which  is  formed  by 
the  interaction  of  hydrogen  sulphide  and  a  lead  compound.  Did  the 
gas  contain  hydrogen  sulphide  ?  Did  the  bottle,  Z),  show  the  formation 
of  carbon  dioxide  ? 

EXERCISES  FOR  THE  CLASS: 

(1)  Describe  briefly  the  whole  experiment. 

(2)  Sketch  the  apparatus. 

(3)  Summarize  the  properties  of  coal  gas. 

Experiment  74.  —  Combustion  of  Illuminating  Gas.  Materials  : 
Pointed  glass  tube  (see  Int.  §  3  (V)),  bottle,  limewater. 

Attach  a  pointed  glass  tube  to  the  rubber  tube  connected  with  the  gas 
jet,  and  lower  a  small  flame  into  a  cold,  dry  bottle.  Observe  at  once 
the  most  definite  result  inside  the  bottle.  Remove  and  extinguish  the 
flame,  add  a  little  limewater  to  the  bottle,  and  shake.  What  are  the 
two  products  of  the  combustion  of  coal  gas  ?  What  do  the  observa- 
tions show  about  the  composition  of  the  main  constituents  of  coal  gas  ? 

Experiment  75.— Construction  of  a  Bunsen  Burner. 

Take  apart  a  Bunsen  burner  and  study  the  construction.  Write 
a  short  description  of  the  burner.  Sketch  the  essential  parts. 


Carbon. 


507 


Experiment  76. —  Bunsen  Burner  Flame.  Materials:  Glass  tube, 
powdered  wood  charcoal,  pin,  copper  wire,  wire  gauze. 

I.  (a)  Close  the  holes  at  the  bottom  of  a  Bunsen  burner  and  hold  a 
glass  tube  in  the  upper  part  of  the  flame.  Note  the  black  deposit. 
What  is  it?  Where  did  it  come  from?  Open  the  holes  and  hold  the 
'blackened  tube  in  the  colorless  flame.  What  becomes  of  the  deposit? 
How  is  the  flame  changed,  if  at  all?  What  does  the  experiment 
suggest  about  the  luminosity  of  flame  ? 

(b)  Dip  a  glass  tube  a  short  distance  into  powdered  wood  charcoal, 
place  the  end  containing  the  charcoal  in  one  of  the  holes  at  the  bottom 
of  the  burner,  and  blow  gently  two  or  three  times  into  the  other  end. 
Describe  and  explain  the  result.  Does  it  verify  the  answer  to  the  last 
question  in  (a)  ? 

(^)  Open  and  close  the  holes  of  a  lighted  burner  several  times. 
Describe  the  result.  Pinch  the  rubber  tube  to  extinguish  the  flame, 
then  light  the  gas  at  the  holes.  What  change  is  produced  in  the  flame? 
What  causes  the  change? 

ANSWER  : 

(1)  What  is  the  object  of  the  holes? 

(2)  Why  does  the  gas  burn  at  the  top  and  not  inside  of  the  burner? 

(3)  Why  does  the  flame  sometimes  "strike  back"  and  burn  inside? 

(4)  Why  is  the  Bunsen  flame  nonluminous? 


II.    (a)   Hold  a  match  across  the  top  of  the  tube  of 
a  lighted  Bunsen  burner.     When  it   begins  to  burn, 
Note  where  it  is  charred, 
and  explain  the  result, 
a  piece  of  wire 
down  upon  the 


remove  and  extinguish  it. 


JL 


Press 

gauze 

flame.      Describe    the 

appearance  of  the  gauze. 

The  same  fact  may  be 
shown  by  sticking  a  pin  through  a  (sul- 
phur) match,  suspending  it  across  the 
burner,  and  then  lighting  the  gas.  The 
position  of  the  match  is  shown  in  Figure 

log.     Turn  on  a  full 

FIG.  no.  — Bent  tube  for  ex-  .  ,     . 

amining  the  structure  of  a  Bun-      current  of  gas  before 
sen  flame.  lighting    it.     What    does    the    whole 


FIG.  109.— Sul- 
phur match  sus- 
pended across  the 
top  of  a  Bunsen 
burner. 


508  Experiments. 

experiment  show  about  the  structure  of  the  lower  part  of  the  Bunsen 
flame?     Verify  your  answer  by  (b). 

(&)  Bend  a  glass  tube  about  15  centimeters  (6  inches)  long  into 
the  shape  shown  in  Figure  no.  Hold  the  shorter  arm  in  the  flame 
about  2  centimeters  (i  inch)  from  the  top  of  the  burner  tube.  Hold 
a  lighted  match  for  an  instant  at  the  upper  end  of  the  tube.  What' 
does  the  result  show  about  the  structure  of^he  Bunsen  flame?  Does  it 
verify  (a)  ? 

(c)  Find  the  hottest  part  of  the  flame,  when  a  full  current  of  gas  is 
burning,  by  holding  a  copper  wire  in  the  flame.  Measure  its  distance, 
approximately,  from  the  top  of  the  burner  tube. 

(d?)  Examine  a  typical  Bunsen  flame  —  one  which  shows  clearly  the 
outlines  of  the  inner  part.  What  is  the  general  shape  of  each  main 
part?  Draw  a  vertical  and  a  cross  section  of  the  flame. 

Experiment  77.  —  Candle  Flame.  Materials :  Candle,  two  blocks 
of  wood,  bottle,  piece  of  stiff  white  paper,  limewater,  matches,  lamp 
chimney,  copper  wire  (  15  centimeters  or  6  inches  long). 

Attach  a  candle  to  a  block  of  wood  by  means  of  a  little  melted  candle 
wax,  and  proceed  as  follows  : — 

(a)  Hold  a  cold,  dry  bottle  over  the  lighted  candle.  Describe  the 
result  produced  inside  the  bottle.  What  is  the  product?  What  is  its 
source?  Remove  the  bottle,  pour  a  little  limewater  into  it,  and  shake. 
Describe  and  explain  the  result.  What  are  the  two  main  products  of  a 
burning  candle? 

(b}  Blow  out  the  candle  flame,  and  immediately  hold  a  lighted  match 
in  the  escaping  smoke.  Does  the  candle  relight?  Why?  What  is  the 
general  nature  of  this  smoke?  How  is  it  related  to  the  candle  wax? 
How  does  (b}  contribute  to  the  explanation  of  (a)  ? 

(c)  Press  a  piece  of  stiff  white  paper  for  an  instant  down  upon  the 
candle  flame  almost  to  the  wick.     Repeat  several  times  with  different 
parts  of  the  paper.      What  does  the  paper  show  about  the  structure  of 
the  flame? 

(d)  Stand  a  lamp  chimney  over  the  lighted  candle.     How  is  the 
flame  effected?     Hold  the  chimney  a  short  distance  (i  centimeter  or 
.5  inch)  above  the  block.     Does  the  candle  continue  to  burn?     Why? 
Keep  the  chimney  in  the  same  position  and  cover  the  top  with  a  block 
of  wood.     What  is  the  result  ?     Why  ? 


Carbon.  509 

(V)  Roll  one  end  of  the  copper  wire  around  a  lead  pencil  to  form  a 
spiral  about  (2  centimeters  or  I  inch)  long.  Press  the  spiral  down 
upon  the  candle  flame.  What  is  the  result?  Why  ? 

EXERCISES  : 

(1)  Draw  a  candle  flame,  showing  the  parts. 

(2)  What  is  the  essential  difference  between  a  candle  flame  and  a 
Bunsen  flame  ? 

(3)  Is  there  any  essential  difference  between  a  candle  flame  and  a 
gas  or  a  lamp  flame  ? 

(4)  Why  do  candles  and  lamps  often  smoke? 

Experiment  78.  —  Kindling  Temperature. 

(«)  Press  a  wire  gauze  down  upon  a  Bunsen  flame.  Where  is  the 
flame  ?  Hold  a  lighted  match  just  above  the  gauze.  Now  where  is 
the  flame  ? 

(<$)  Extinguish  the  flame.  Turn  on  the  gas,  hold  the  gauze  in  the 
escaping  gas,  about  5  centimeters  (2  inches)  above  the  top  of  the  burner, 
and  thrust  a  lighted  match  into  the  gas  above  the  gauze.  Where  is  the 
flame  ?  Lower  the  gauze  slowly  and  describe  the  final  result. 

(c}  Hold  the  gauze  in  the  flame  in  one  position  for  a  minute  or  two. 
Where  is  the  flame  at  the  end  of  this  time  ?  Why  ? 

EXERCISES  : 

(1)  Define  kindling  temperature. 

(2)  What  application  is    made  of  the  principle  illustrated    by  this 
experiment  ? 

(3)  State  exactly  how  this  experiment  illustrates  kindling  tempera- 
ture. 

Experiment  79. — Reduction  and  Oxidation  with  the  Blow- 
pipe. Materials :  Blowpipe,  blowpipe  tube,  charcoal,  lead  oxide 
(litharge),  sodium  carbonate,  sodium  sulphate,  wood  charcoal,  silver 
coin,  zinc,  lead,  tin. 

Slip  the  blowpipe  tube  into  the  burner,  light  the  gas  and  lower  the 
flame  until  it  is  about  4  centimeters  (1.5  inches)  high.  Rest  the  tip  of  the 
blowpipe  on  the  top  of  the  tube,  placing  the  tip  just  within  the  flame. 
Put  the  other  end  of  the  blowpipe  between  the  lips,  puff  out  the  cheeks, 
inhale  through  the  nose,  and  exhale  into  the  tube,  using  the  cheeks  some- 
what as  a  bellows.  Do  not  blow  in  puffs,  but  produce  a  continuous  flow 
of  air  by  steady  and  easy  inhaling  and  exhaling.  The  operation  is  nat- 


Experiments. 

ural  and  simple,  and,  if  properly  performed,  will  not  make  one  out  of 
breath.  The  flame  should  be  an  inner  blue  cone  surrounded  by  an  outer 
and  almost  invisible  cone,  though  its  shape  varies  with  the  method  of 
production  (see  Fig.  44).  Practice  until  the  flame  is  produced  volun- 
tarily and  without  exhaustion.  Watch  the  flame  and  learn  to  distin- 
guish the  two  parts,  so  that  they  may  be  intelligently  utilized. 

I.  Reduction,     (a)  Make  a  shallow  hole  at  one  end  of  the  flat  side  of 
a  piece  of  charcoal.     Fill  the  hole  with  a  mixture  of  equal  parts  of  pow- 
dered sodium  carbonate  and  lead  oxide,  and  heat  the  mixture  in  the 
reducing  flame.     The  sodium  carbonate  melts   and  assists  the  fusion 
of  the  oxide,  but  the  former  is  not  changed  chemically.     In  a  short  time 
bright,  silvery  globules  will  appear  on  the  charcoal.     Let  the  mass  cool, 
and  pick  out  the  largest  globules.    Put  one  or  two  in  a  mortar,  and  strike 
with  a  pestle.    Are  they  soft  and  malleable,  or  brittle  and  hard  ?     State 
the  result  when  a  globule  is  drawn  across  or  rubbed  upon  a  white  paper. 
How  do  the  properties  compare  with  those  of  metallic  lead  ?    What  has 
become  of  the  oxygen  ?     Of  what  chemical  use  is  the  charcoal  ? 

(b)  Grind  together  in  a  mortar  a  little  sodium  sulphate  and  wood 
charcoal,  adding  at  intervals  just  enough  water  to  hold  the  mass  to- 
gether. Heat  this  paste  fora  few  minutes  in  the  reducing  flame  as  in 
(a) .  Scrape  the  fused  mass  into  a  test  tube,  boil  in  a  little  water,  and 
put  a  drop  of  the  solution  on  a  bright  silver  coin.  If  a  dark  brown  stain 
is  produced,  it  is  evidence  of  the  formation  of  silver  sulphide.  Repeat, 
if  no  such  stain  is  produced.  State  all  the  chemical  changes  which  led 
to  the  production  of  the  silver  sulphide,  explaining  at  the  same  time 
how  the  experiment  illustrates  reduction. 

II.  Oxidation,     (a)  Heat  a  small  piece  of  zinc  on  charcoal  in  the 
oxidizing  flame.     What  is  the  product  ?     Observe  its  color,  and  the  color 
of  the  coating  on  the  charcoal  when  hot  and  cold.     Record  as  described 
\&(d). 

(b)  Heat  a  piece  of  lead  as  in  (a}.     Observe  the  presence  or  absence 
of  fumes,  as  well  as  the  color  of  the  coating  when  hot  and  cold.     See  (d). 

(c)  Heat  a  small  piece  of  tin  in  the  oxidizing  flame.    Observe  as  in  (b} . 
(d}  Tabulate   the  above  observations,  stating  (i)  the  color  of  the 

hot  and  cold  coating  on  the  charcoal,  (2)  presence  or  absence  of  fumes, 
(3)  name  of  product. 

EXERCISES  : 

(1)  Sketch  a  blowpipe. 

(2)  Sketch  a  flame  showing  the  oxidizing  and  reducing  parts. 


Fluorine,   Bromine,  and  Iodine.  511 


FLUORINE,  BROMINE,  AND   IODINE. 

Experiment  80.  —  Preparation  and  Properties  of  Hydrofluoric 

Acid.    Materials :  Lead  dish,  glass  plate,  paraffin,  file,  calcium  fluoride, 
concentrated  sulphuric  acid. 

Precaution.  Hydrofluoric  acid  gas  is  a  corrosive  poison.  An  aque- 
ous solution  of  the  gas  —  commercial  hydrofluoric  acid —  burns  the  flesh 
frightfully. 

Warm  a  glass  plate  about  10  centimeters  (4  inches)  square  by  dipping 
it  into  hot  water  or  by  standing  it  near  a  warm  object,  such  as  a  radiator. 
If  it  is  held  over  a  flame,  it  is  liable  to  crack.  Coat  one  surface  with 
paraffin.  The  surface  should  be  uniformly  covered  with  a  thin  layer. 
Scratch  letters,  figures,  or  a  diagram  through  the  wax  with  a  file.  Be 
sure  the  instrument  removes  the  wax  through  to  the  glass,  and  that  the 
lines  are  not  too  fine. 

Put  5  grams  of  calcium  fluoride  in  a  lead  dish  and  add  just  enough 
concentrated  sulphuric  acid  to  form  a  thin  paste.  Stir  the  mixture  with 
a  file.  Place  the  glass  plate,  wax  side  down,  upon  the  lead  dish  and 
stand  the  whole  apparatus  in  the  hood  for  several  hours,  or  until  some 
convenient  time.  Remove  the  plate.  Scrape  the  contents  of  the  dish, 
immediately,  into  a  waste  jar  in  the  hood,  and  wash  the  dish  free  from 
acid.  Most  of  the  wax  can  be  scraped  from  the  glass  plate  with  a  knife. 
The  last  portions  can  be  removed  by  rubbing  with  a  cloth  moistened 
with  alcohol  or  turpentine.  Do  not  attempt  to  melt  off 
the  wax  over  the  flame.  If  the  experiment  has  been 
properly  performed,  the  plate  will  be  etched  where  the 
glass  was  exposed  to  the  hydrofluoric  acid  gas. 


Experiment  81.  —  Preparation  and  Properties  of 
Bromine.  Materials :  Potassium  bromide,  manganese 
dioxide,  dilute  sulphuric  acid,  bottle  of  water,  test-tube 
holder.  The  apparatus  is  shown  in  Figure  1 1 1 .  The 
large  test  tube  is  provided  with  a  one-hole  rubber 
stopper  to  which  is  fitted  the  bent  glass  tube.  The 
latter  is  about  30  centimeters  (12  inches)  long,  and  is 
bent  according  to  the  directions  given  in  the  Introduc-  paratus  for  pre. 
tion,  §  3  (b).  paring  bromine. 


512  Experiments. 

Precaution.  Bromine  is  a  corrosive  liquid  which  forms,  at  the 
ordinary  temperature,  a  suffocating  vapor.  Perform  in  the  hood  all 
experiments  which  use  or  evolve  bromine.  - 

Put  a  dozen  crystals  of  potassium  bromide  in  the  test  tube,  add  an 
equal  quantity  of  manganese  dioxide  and  10  cubic  centimeters  of  dilute 
sulphuric  acid.  Insert  the  stopper  and  its  tube  securely,  and  boil 
gently.  Do  not  hold  the  test  tube  in  the  hand,  but  use  the  test  tube 
holder.  Brown  fumes  soon  appear  in  the  test  tube  and  pass  out  of  the 
delivery  tube.  Regulate  the  heating  so  that  this  vapor  will  condense 
and  collect  in  the  lower  bend  of  the  delivery  tube.  Both  vapor  and 
liquid  are  bromine.  When  no  further  boiling  produces  bromine  vapor 
in  the  test  tube,  pour  the  bromine  from  the  delivery  tube  into  a  bottle 
of  water.  Observe  and  record  the  physical  properties  of  this  bromine, 
especially  the  color,  solubility  in  water,  specific  gravity,  volatility,  and 
physical  state.  Try  the  action  of  the  contents  of  the  bottle  on  litmus 
paper ;  if  the  action  is  not  marked,  push  the  paper  down  near  the  bro- 
mine. Determine  the  odor  by  smelling  cautiously  of  the  water  in  the 
bottle.  As  soon  as  these  observations  have  been  made,  pour  the  con- 
tents of  the  bottle  into  the  sink  and  flush  with  water,  or  pour  into  a  jar 
in  the  hood.  Wash  the  test  tube  free  from  all  traces  of  bromine,  taking 
care  to  get  none  on  the  hands. 

ANSWER  : 

(1)  In  what  ways  does  bromine  physically  resemble  chlorine?     In 
what  ways  does  it  differ  from  chlorine? 

(2)  How  is  it  essentially  different  from  all  other  elements  previously 
studied  ? 

Experiment  82.  —  Properties  of  Potassium  Bromide.  Materials : 
Potassium  bromide,  silver  nitrate  solution,  ammonium  hydroxide. 

Examine  a  crystal  of  potassium  bromide,  and  state  its  most  obvious 
properties.  Dissolve  it  in  a  test  tube  half  full  of  water,  and  add  a  few 
drops  of  silver  nitrate  solution.  Describe  the  result.  Is  the  solid  prod- 
uct soluble  in  ammonium  hydroxide?  How  can  bromides  be  distin- 
guished from  chlorides  ?  Do  the  properties  of  bromides,  typified  by 
potassium  bromide,  suggest  any  marked  relation  to  chlorides  ? 

Experiment  83.  —  Preparation  and  Properties  of  Iodine.  Ma- 
terials: Potassium  iodide,  manganese  dioxide,  mortar  and  pestle,  con- 
centrated sulphuric  acid,  funnel,  cotton. 


Fluorine,  Bromine,  and  Iodine.  513 

Grind  together  in  a  mortar  a  dozen  large  crystals  of  potassium  iodide 
and  about  twice  the  bulk  of  manganese  dioxide.  Put  the  mixture  in  a 
test  tube  provided  with  a  holder,  moisten  with  water,  and  add  a  few 
cubic  centimeters  of  concentrated  sulphuric  acid.  Plug  with  cotton  the 
inside  opening  of  a  funnel,  and  hold  the  latter  firmly  over  the  mouth  of 
the  test  tube.  Heat  the  test  tube  gently  with  a  low  flame  (5  centime- 
ters or  2  inches).  The  vapor  of  iodine  will  fill  the  test  tube,  and  crys- 
tals will  collect  in  the  upper  part  of  the  test  tube  and  in  the  funnel. 
If  the  crystals  collect  in  the  test  tube,  a  gentle  heat  will  force  them 
into  the  funnel.  Continue  to  heat  until  enough  iodine  collects  in  the 
funnel  for  several  experiments.  Scrape  the  crystals  into  a  dish. 

Study  the  properties  as  follows  :  — 

(«)  Observe  and  record  the  physical  properties  of  iodine,  especially 
the  color  of  the  solid  and  of  the  vapor,  volatility,  and  odor  (cautiously). 

(£)  Heat  a  crystal  in  a  dry  test  tube,  and  when  the  tube  is  half  full 
of  vapor,  invert  it.  What  does  the  result  show  about  the  density  of 
iodine  vapor? 

(c)  Touch  a  crystal  with  the  finger.  What  color  is  the  stain  ?  Will 
water  remove  it?  Will  alcohol?  Will  a  solution  of  potassium  iodide? 
What  do  these  results  show  about  the  solubility  of  iodine  ? 

(NOTE.  —  If  crystals  are  left,  use  them  in  the  next  experiment.  Pre- 
serve in  a  stoppered  bottle.) 

Experiment  84.  —  Test  for  Iodine  with  Carbon  Bisulphide. 
Materials :  Iodine,  potassium  iodide,  carbon  disulphide,  chlorine  water. 

Precaution.  Carbon  disulphide  is  inflammable.  It  should  not  be 
used  near  flames. 

(a)  Free  iodine.     Add  a  few  drops  of  carbon  disulphide  to  a  very 
dilute  solution  of  iodine,  made  by  dissolving  a  crystal  of  iodine  in  a 
solution  of  potassium  iodide,  and  observe  the  color  of  the  carbon  disul- 
phide, which,  being  much  heavier  than  water,  will  sink  to  the  bottom  of 
the  test  tube.     How  does  it  resemble  the  color  of  iodine  vapor  ? 

(b)  Combined  iodine.     Add  a  few  drops  of  carbon  disulphide  to  a 
very  dilute  solution  of  potassium  iodide.     Is  there  positive  evidence  of 
iodine  ?     Now  add  several  drops  of  chlorine  water,  and  shake.     How 
does  this  result  compare  with  the  final  result  in  (a)  ?     The  result  is  due 
to  the  fact  that  chlorine  liberates  iodine  from  its  compounds,  and  the 
iodine,  being  free,  exhibits  the  characteristic  color. 


Experiments. 

Experiment  85.  —  Test  for  Iodine  with  Starch.  Materials: 
Starch,  mortar  and  pestle,  iodine  solution,  potassium  iodide,  chlorine 
water. 

Grind  a  lump  of  starch  in  a  mortar  with  a  little  water  to  creamy  con- 
sistency. Pour  this  into  about  100  cubic  centimeters  of  boiling  water, 
and  stir  the  hot  liquid.  Allow  it  to  cool,  or  cool  it  by  holding  the 
vessel  in  a  stream  of  cold  water,  and  then  pour  off  the  clear  liquid. 
Use  this  cold  starch  solution  to  test  for  iodine. 

(a)  Free  iodine.     Add  a  few  cubic  centimeters  of  the  starch  solution 
to  a  test  tube  nearly  full  of  water,  and  then  add  a  few  drops  of  iodine 
solution.     The  deep  blue  color  is  due  to  the  presence  of  a  compound 
which  is  always  formed  under  these  circumstances,  but  the  composition 
of  which  is  unknown.     If  the  color  is  black,  pour  out  half  of  the  liquid 
and  add  more  water,  or  pour  some  of  the  liquid  into  a  dish  of  water. 

(b)  Combined  iodine.     Add  a  few  cubic  centimeters  of  the  starch 
solution  to  a  very  dilute  solution  of  potassium  iodide.     Is  the  blue  com- 
pound formed  ?    Add  a  few  drops  of  chlorine  water,  and  shake.     Com- 
pare with  the  final  result  in  Experiment  84  (b) . 

Experiment  86.  —  Detection  of  Starch  by  Iodine.  Materials: 
Dilute  solution  of  iodine  (in  potassium  iodide),  mortar  and  pestle, 
potato,  rice,  bread. 

Test  the  potato,  rice,  and  bread  for  starch  by  grinding  a  little  of  each 
with  water  in  a  mortar,  and  then  adding  a  few  drops  of  the  extract  to  a 
very  dilute  solution  of  iodine.  State  the  result  in  each  case. 

Experiment  87.  —  Properties  of  Potassium  Iodide.  Materials : 
Potassium  iodide,  silver  nitrate  solution,  ammonium  hydroxide. 

Proceed  with  the  potassium  iodide  as  in  Experiment  82. 

How  can  iodides  be  distinguished  from  chlorides  ?  Do  iodides, 
typified  by  potassium  iodide,  suggest  any  marked  relation  to  bromides 
and  chlorides  ? 

SULPHUR  AND  ITS  COMPOUNDS. 

Experiment  88. — Properties  of  Sulphur. 

(a)  Examine  a  lump  of  sulphur,  and  state  briefly  its  most  obvious 
physical  properties. 

(6)  Optional.  Weigh  a  lump  of  roll  sulphur  to  a  decigram.  Slip  it 
carefully  into  a  graduated  cylinder  previously  filled  with  water  to  a 


Sulphur  and  its  Compounds.  515 

known  point  —  about  half  full  —  and  note  the  increase  in  the  volume  of 
water.  This  increase  in  volume  is  equal  to  the  volume  of  the  sulphur. 
Calculate  the  specific  gravity  of  sulphur  from  the  observed  data. 

(NOTE.  —  Specific  gravity  equals  weight  in  air  divided  by  weight  of 
equal  volume  of  water.) 

Experiment  89.  —  Amorphous  Sulphur.  Materials:  Sulphur,  old 
test  tube,  evaporating  dish. 

Put  a  few  pieces  of  roll  sulphur  in  an  old  test  tube.  Heat  carefully 
until  the  sulphur  boils,  and  then  quickly  pour  the  contents  of  the  test 
tube  into  a  dish  of  cold  water.  This  is  amorphous  sulphur.  Note  its 
properties.  Preserve,  and  examine  it  after  twenty -four  hours.  Describe 
the  change,  if  any. 

Define  amorphous,  and  illustrate  it  by  this  experiment. 

Experiment  90.  —  Crystallized  Sulphur. J  Materials  :  Sulphur 
(roll  and  flowers),  Hessian  crucible,  carbon  disulphide,  evaporating 
dish. 

(a)  Monodinic.  Fill  a  small  Hessian  crucible  nearly  full  of  roll 
sulphur.  Support  the  crucible  in  the  ring  of  an  iron  stand,  and  heat 
until  all  the  sulphur  is  melted.  Let  it  cool,  and  as  soon  as  crystals 
shoot  out  from  the  walls  just  below  the  surface,  pour  the  remaining 
melted  sulphur  into  a  dish  of  cold  water.  When  the  crucible  can  be 
handled  without  discomfort,  crack  it  open  lengthwise.  Observe  and 
record  the  properties  of  the  crystals,  especially  the  shape,  size,  color, 
luster,  brittleness,  and  any  other  characteristic  property.  Allow  the 
best  crystals  to  remain  undisturbed  for  a  day  or  two ;  then  examine 
again,  and  record  any  marked  changes. 

(b}  Orthorhombic.  Put  3  grams  of  flowers  of  sulphur  in  a  test  tube 
and  add  about  5  cubic  centimeters  of  carbon  disulphide  —  remember 
the  precaution  to  be  observed  in  using  this  liquid  (see  Exp.  84). 
Shake  until  all  the  sulphur  is  dissolved,  then  pour  the  clear  solution 
into  an  evaporating  dish  to  crystallize.  It  is  advisable,  though  not 
absolutely  necessary,  to  stand  the  dish  in  the  hood  or  out  of  doors, 
where  there  is  no  flame  and  where  the  offensive  vapor  will  be  quickly 
removed.  Watch  the  crystallization  toward  the  end,  and,  if  perfect 
crystals  form,  remove  them  with  the  forceps  (see  Fig.  49).  Allow  the 

i  See  Appendix,  §  3  (3),  (5). 


Experiments. 

liquid  to  evaporate  almost  entirely,  then  remove  and  dry  the  crystals. 
Examine  them  as  in  (a)  and  record  their  properties. 
EXERCISES  : 

(1)  Tabulate  the  essential  results  in  (a)  and  (£). 

(2)  Make  an  outline  sketch  of  an  orthorhombic  crystal  of  sulphur. 

Experiment  91.  —  Combining  Power  of  Sulphur.  Materials  : 
Sulphur,  deflagrating  spoon,  bottle,  iron  powder,  hydrochloric  acid. 

(a)  Set  fire  to  a  little  sulphur  in  a  deflagrating  spoon,  and  lower 
the  spoon  into  a  bottle.  Cautiously  waft  the  fumes  toward  the  nose, 
and  observe  and  describe  the  odor.  The  product  is  a  mixture  of  two 
oxides  of  sulphur.  What  does  their  formation  show  about  the  combin- 
ing power  of  sulphur  ? 

(£)    Repeat  Experiment  34. 

Results  similar  to  that  in  (<£)  are  obtained  with  copper  and  other 
metals.  Draw  a  general  conclusion  regarding  the  power  of  sulphur  to 
combine  with  metals. 

Experiment  92.  —  Sulphur  and  Matches. 

(a)  Examine  a  sulphur  match.    Do  you  detect  any  sulphur  ?    Where? 

(b)  Light  a  sulphur  match,  and  observe  the  entire  action,  as  far  as 
the  sulphur  is  concerned.     Describe  it. 

(c)  What  is  the  function  of  the  sulphur  in  a  burning  match  ? 

Experiment  93.  —  Preparation  of  Hydrogen  Sulphide.  Mate- 
rials: Ferrous  sulphide,  dilute  hydrochloric  acid,  three  bottles,  three 
glass  plates,  stoppered  bottle,  litmus  paper.  Use  the  same  apparatus 
as  in  Experiment  38. 

Precaution.  Hydrogen  sulphide  is  a  poisonous  gas  and  has  an 
offensive  odor.  It  should  not  be  inhaled.  Perform  in  the  hood  all  ex- 
periments evolving  hydrogen  sulphide. 

(a)  Construct  and  arrange  an  apparatus  like  that  shown  in  Figure  104. 
Fill  the  test  tube,  A,  one  third  full  of  coarsely  powdered  ferrous  sulphide, 
insert  the  stopper  tightly,  pour  enough  hydrochloric  acid  through  the 
safety  tube  to  cover  the  contents  of  the  test  tube.  Hydrogen  sulphide 
gas  is  rapidly  evolved.  If  the  evolution  of  gas  slackens  or  stops,  warm 
gently  or  add  more  hydrochloric  acid.  Collect  three  bottles,  removing 
each  as  soon  as  full  and  covering  with  a  glass  plate.  Set  aside  until 
needed. 


Sulphur  and  its  Compounds.  517 

(£)  As  soon  as  the  last  bottle  of  gas  has  been  removed  and  covered, 
put  in  its  place  a  bottle  one  fourth  full  of  water.  Adjust  its  height  (by 
wooden  blocks  or  by  lowering  the  generator)  so  that  the  end  of  the 
delivery  tube  reaches  to  the  bottom  of  the  bottle.  Continue  to  pass 
the  gas  into  the  water,  by  heating  the  test  tube  if  necessary.  The  gas 
will  be  absorbed  by  the  water,  forming  hydrogen  sulphide  water. 
Preserve  it  in  a  stoppered  bottle  for  Experiment  95. 

Proceed  at  once  with  next  experiment. 

Experiment  94.  —  Properties  of  Hydrogen  Sulphide  Gas. 

Study  as  follows  the  hydrogen  sulphide  gas  prepared  in  Experi- 
ment 93  :  — 

(#)  Waft  a  little  of  the  gas  cautiously  toward  the  nose,  and  describe 
the  odor.  This  is  characteristic  of  hydrogen  sulphide,  and  is  a  decisive 
test.  Has  the  gas  color? 

(b)  Test  the  gas  from  the  same  bottle  with  both  kinds  of  moist 
litmus  paper.     Is  it  acid,  alkaline,  or  neutral  ? 

(c)  Bring  a  lighted  match  to  the  mouth  of  the  same  bottle.     Observe 
the  properties  of  the  flame  as  in  previous  experiments.     Observe  cau- 
tiously the  odor  of  the  product  of  the  burned  gas ;  to  what  compound 
is  the  odor  due?     What,  then,  is  one  component  of  hydrogen  sulphide  ? 

(d)  Burn  another  bottle  of  hydrogen  sulphide  and  hold  a  cold  bottle 
over  the  burning  gas.     What  additional  experimental  evidence  does  this 
result  give  regarding  the  composition  of  hydrogen  sulphide  ? 

(V)  Repeat  any  of  the  above  with  the  remaining  bottle  of  gas. 

EXERCISES  : 

(1)  Summarize  the  properties  of  hydrogen  sulphide  gas. 

(2)  State  the  experimental  evidence  of  its  composition. 

Experiment  95.  —  Preparations  and  Properties  of  some  Sul- 
phides. Materials:  Hydrogen  sulphide  water  prepared  in  Experiment 
93,  clean  copper  wire,  clean  sheet  lead,  bright  silver  coin,  lead  oxide 
(litharge)  ;  solutions  of  lead  nitrate,  arsenic  trioxide  (in  hydrochloric 
acid),  tartar  emetic,  zinc  sulphate. 

(a)  Shake  the  bottle  of  hydrogen  sulphide  water  prepared  in  Experi- 
ment 93  (or  a  similar  solution),  and  hold  successively  at  the  mouth  or 
in  the  neck  of  the  bottle  (i)  a  clean  copper  wire,  (2)  a  bright  strip 
of  lead,  and  (3)  an  untarnished  silver  coin.  Describe  the  result  in 
each  case.  These  compounds  are  sulphides  of  the  respective  metals. 


518  Experiments. 

(£)  Put  a  little  litharge  —  the  brownish  yellow  oxide  of  lead  — 
in  a  test  tube,  cover  it  with  hydrogen  sulphide  water,  and  warm 
gently.  The  product  is  lead  sulphide.  Describe  it.  Explain  the 
change. 

(c)  Add  hydrogen  sulphide  water  to  lead  nitrate  solution.  The 
product  is  lead  sulphide.  Observe  the  color. 

(*/)  Proceed  as  in  (c)  with  the  arsenic  solution.  Observe  the  color 
of  the  arsenic  sulphide. 

(e)  Proceed  as  in  (c)  with  the  tartar  emetic  solution.  Tartar  emetic 
is  a  compound  of  antimony.  Observe  the  color  of  the  antimony 
sulphide. 

(_/")  Proceed  as  in  (c)  with  the  zinc  sulphate  solution.  Observe  the 
color  of  the  zinc  sulphide. 

Experiment  96.  — Preparation  of  Sulphur  Dioxide.  Materials  : 
Sodium  sulphite,  concentrated  sulphuric  acid,  litmus  paper,  three 
bottles,  two  glass  plates,  stick  of  wood,  pink  flower.  The  apparatus  is 
constructed,  arranged,  and  used  as  in  Experiment  41,  with  one  excep- 
tion. The  safety  tube  must  be  replaced  by  a  'dropping  tube  made  thus . 
Cut  off  the  top  of  a  thistle  tube  about  2.5  centimeters  (i  inch)  below 
the  juncture  of  the  stem  and  cup,  slip  a  short  rubber  tube  (5  centi- 
meters, 2  inches,  long  )  over  one  end  of  the  stem,  attach  a  Mohr?s  pinch- 
cock  to  the  rubber  tube,  and  connect  the  tube  with  the  cup. 

(a)  Put  about  10  grams  of  sodium  sulphite  in  the  large  test  tube, 
cover  with  water,  and  insert  the  stopper  with  its  tubes.     Adjust  the  ap- 
paratus as  shown  in  Figure  104.     Fill  the  cup  with  concentrated  sulphuric 
acid,  open  the  pinchcock  a  little,  and  let  the  acid  flow  drop  by  drop  upon 
the  sodium  sulphide.     Sulphur  dioxide  gas  is  evolved  and  passes  into 
the  bottle,  which  should  be  removed  when  full,  as  previously  described. 
Moist  blue  litmus  paper  held  at  the  mouth  of  the  bottle  will  show  when 
the  latter  is  full.     Collect  two  bottles  of  gas,  cover  each  with  a  glass 
plate,  and  set  aside  until  needed. 

(b)  As  soon  as   the  second   bottle  of  gas  has  been  removed  and 
covered,  put  in  its  place  a  bottle  one  fourth  full  of  water.     Adjust  its 
height  (if  necessary)  by  wooden  blocks,  so  that  the  end  of  the  delivery 
tube  is  just  above  the  surface.     Continue  to  add  the  acid  drop  by  drop, 
at  intervals,  and  the  gas  will  be  absorbed  by  the  water.     Shake  the  bottle 
occasionally. 

Meanwhile  study  the  gas  already  collected. 


Sulphur  and  its  Compounds.  519 

Experiment  97. — Properties  of  Sulphur  Dioxide  Gas. 

Proceed  as  follows  with  the  gas  prepared  in  Experiment  96  (#)  :  — 

(a)  Observe   and  state   the  most  obvious  physical  properties,  e.g. 
color,  odor  (cautiously),  density. 

(b)  Hold  a  blazing  stick  in  a  bottle  of  the  gas.     Will  the  gas  burn 
or  support  combustion  ?     What  previously  acquired  facts  would  have 
enabled  you  to  predict  this  result  ? 

(c)  Pour  water  into  the  same  bottle  of  sulphur  dioxide  until  half  full, 
cover  with  the  hand,  and  shake.     What  is  the  evidence  of  solution  ? 
Is  the  resulting  liquid  acid,  alkaline,  or  neutral  ? 

(d)  Moisten  a  pink  flower  with  a  few  drops  of  water,  hang  it  in  the 
remaining  bottle  of  sulphur  dioxide,  holding  it  in  place  by  putting  the 
stem  between  the  glass  and  a  cork.     Observe  and  describe  any  change 
in  the  color  of  the  flower.     What  is  this  operation  called  ? 

Experiment  98. —  Properties  of  Sulphurous  Acid. 

Test  as  follows  the  solution  of  sulphurous  acid  prepared  in  Experi- 
ment 96  (£)  :  — 

(a)  Taste  cautiously,  and  describe  the  result. 

(b)  Apply  the  litmus  test,  and  state  the  result. 

(c}  Pour  a  few  drops  of  concentrated  sulphuric  acid  into  the  bottle. 
What  gas  is  liberated  ? 

Experiment  99. — Action  of  Sulphuric  Acid  with  Organic  Matter. 

Materials :  Concentrated  sulphuric  acid,  sheet  of  white  paper,  sugar, 
starch,  stick  of  wood. 

(a)  Write  some  letters  or  figures  with  dilute  sulphuric  acid  on  a  sheet 
of  white  paper,  and  move  the  paper  back  and  forth  over  a  low  flame,  taking 
care  not  to  set  fire  to  the  paper.  As  the  water  evaporates  the  dilute 
acid  becomes  concentrated.  Observe  and  describe-the  result.  Paper  is 
largely  a  compound  of  carbon,  hydrogen,  and  oxygen,  and  the  hydrogen 
and  oxygen  are  present  in  the  proportion  to  form  water.  Explain  the 
general  chemical  change  in  this  experiment. 

(b}  Fill  a  test  tube  one  fourth  full  of  sugar,  add  an  equal  bulk  of 
water,  stand  the  test  tube  in  the  rack,  and  add  cautiously  several  drops 
of  concentrated  sulphuric  acid.  If  there  is  no  decided  result,  add 
more  acid.  What  is  the  black  product  ?  Compare  the  final  result  with 
that  obtained  in  Experiment  59  (£).  Is  the  chemical  action  the  same  in 


520  Experiments. 

each  experiment  ?  Are  the  statements  made  in  (a)  about  paper  also 
true  of  sugar  ? 

(c)  Repeat  (£),  using  powdered  starch  instead  of  sugar.  Describe 
the  result.  How  does  the  result  resemble  that  in  (b}  and  in  Experi- 
ment 59  (a)  ?  Predict  the  components  of  starch.  In  what  simple  way 
may  the  prediction  be  verified  ? 

(//)  Stand  a  stick  of  wood  in  a  test  tube  one  fourth  full  of  concen- 
trated sulphuric  acid.  Allow  it  to  remain  in  the  acid  for  fifteen  minutes, 
then  remove  the  stick  and  wash  off  the  acid.  Describe  the  change  in 
the  stick.  Does  it  resemble  that  in  (#),  (£),  and  (c),  and  in  Experiment 
59<X>? 

Experiment   100.  —  Test    for    Sulphuric    Acid  and   Sulphates. 

Materials:  Sulphuric  acid,  sodium  sulphate,  barium  chloride  solution, 
calcium  sulphate,  charcoal,  powdered  charcoal,  blowpipe,  silver  coin. 

(a)  Repeat  Experiment  12  (c)  with  sulphuric  acid  and  with  sodium 
sulphate  solution. 

(£)  Repeat  Experiment  79  I  (£)  with  calcium  sulphate  instead  of 
sodium  sulphate. 

EXERCISES  : 

(1)  State  briefly  the  test  for  sulphuric  acid  and  soluble  sulphates. 
For  insoluble  sulphates. 

(2)  How  can  a  sulphate  be  distinguished  from  a  sulphite  ? 

SILICON  AND  BORON. 

Experiment  101.  —  Preparation  and  Properties  of  Silicic  Acid. 

Materials:  Sodium  silicate  solution,  hydrochloric  acid,  evaporating 
dish,  gauze-covered  ring. 

Add  dilute  hydrochloric  acid  to  a  test  tube  half  full  of  sodium  silicate 
solution,  and  shake.  The  jellylike  precipitate  is  silicic  acid.  Rub 
some  between  the  fingers  and  describe  the  result.  Evaporate  the 
precipitate  to  dryness  in  a  porcelain  dish  which  stands  upon  a  gauze- 
covered  ring  in  the  hood.  As  the  mass  hardens,  stir  it  with  a  glass  rod. 
Toward  the  end,  add  more  hydrochloric  acid  and  evaporate  to  complete 
dryness.  Then  heat  strongly  for  five  minutes.  The  residue  is  silicon 
dioxide  mixed  with  chlorides  of  sodium  and  potassium.  Rub  some 
between  the  fingers  or  across  a  glass  plate.  Is  any  grit  detected  ?  State 
the  chemical  changes  which  occur  in  changing  sodium  silicate  into 
silicon  dioxide. 


Silicon  and  Boron.  521 

Experiment  102. — Tests  with  Borax  Beads.  Materials:  Pow- 
dered borax,  platinum  test  wire  (see  Int.  §  5  (4)),  solutions  of  cobalt 
nitrate  and  copper  sulphate,  manganese  dioxide. 

Make  a  small  loop  on  the  end  of  the  platinum  test  wire,  moisten  it, 
and  dip  it  into  powdered  borax.  Heat  it  in  the  flame,  rotating  it  slowly  ; 
at  first  the  borax  swells,  but  finally  shrinks  to  a  small,  transparent 
bead.  If  the  bead  is  too  small  add  more  borax  and  heat  again.  After 
use,  the  bead  may  be  removed  by  dipping  it,  white  hot,  into  water ;  the 
sudden  cooling  shatters  the  bead,  which  may  then  be  easily  rubbed  or 
scraped  from  the  wire. 

(a)  Cobalt  Compounds.     Touch   a  transparent  borax  bead  with  a 
glass  rod  which  has  a  drop  of  cobalt  nitrate  solution  on  the  end.     Heat 
the  bead  in  the  oxidizing  flame.     Observe  the  color  when  cold.     If  it  is 
black  melt  a  little  more  borax  into  the  bead  ;    if  faintly  colored,  moisten 
again  with  the  cobalt  solution.     The  color  is  readily  detected  by  look- 
ing at  the  bead  against  a  white  object  in  a  strong  light,  or  by  examining 
it  with  a  lens.     When  the  color  has  been  definitely  determined,  heat 
again  in  the  reducing  flame.     Compare  the  color  of  the  cold  bead  with 
the  previous  observation. 

(b)  Copper  compozinds.    Make  another  transparent  bead,  moisten  it 
with  copper  sulphate  solution  and  heat  it  first  in  the  oxidizing  flame, 
and  then  in  the  reducing  flame.   Compare  the  colors  of  the  cold  beads, 
and  draw  a  conclusion. 

(c)  Manganese  Compounds.    Make  another  transparent  bead,  touch 
it  with  a  minute  quantity  of  manganese  dioxide,  and  proceed  as  in  (b). 
Compare  the  colors  of  the  cold  beads,  and  draw  a  conclusion. 

(d)  Tabulate  the  results  of  this  experiment. 

EXERCISE  : 

Draw  a  Bunsen  flame,  showing  the  reducing  and  oxidizing  parts. 

Experiment  103.  —  Preparation  and  Properties  of  Boric  Acid 
and  the  Test  for  Boron.  Materials:  Borax,  alcohol,  evaporating 
dish,  concentrated  hydrochloric  acid. 

To  a  test  tube  half  full  of  boiling  water,  add  about  10  grams  of 
powdered  borax.  Add  about  5  cubic  centimeters  of  concentrated 
hydrochloric  acid  to  this  hot  solution,  and  let  the  whole  cool.  Crystals 
of  boric  acid  will  separate.  Filter.  Describe  the  crystals. 

Put  some  of  the  crystals  in  an  evaporating  dish,  add  a  little  alcohol, 


522  Experiments. 

and  set  fire  to  the  solution.  Observe  the  color  of  the  flame.  It  is 
caused  by  a  complex  compound  of  boron,  and  is  the  test  for  this 
element. 

PHOSPHORUS,  ARSENIC,  ANTIMONY,  AND  BISMUTH. 
Experiment  104.  —  Some  Properties  of  Phosphorus. 

(a)  Smell  of  the  head  of  a  phosphorus-tipped  match.     Describe  the 
odor. 

(b)  Rub  the  head  of  a  phosphorus-tipped  match  in  a  dark  place,  and 
observe  and  describe  the  result. 

(c)  The  most  striking  property  of  phosphorus  is  the  readiness  with 
which  it  lights  and  burns  in  air.     This  property  is  too  dangerous  to  try 
in  the  laboratory.     Read  about  it  in  the  text  book.     What  application 
is  made  of  this  property  ?     Why  ? 

Experiment  105.  —  Test  for  Arsenic. 

Repeat  Experiment  95  (</). 

Experiment  106.  —  Test  for  Antimony. 

Repeat  Experiment  95  (e). 

Experiment  107.  —  Test  for  Bismuth.  Materials:  Bismuth,  aqua 
regia. 

Prepare  a  solution  of  bismuth  chloride  by  heating  the  metal  with  aqua 
regia.  Fill  the  test  tube  half  full  of  water,  and  describe  the  result.  The 
product  is  bismuth  oxychloride.  How  is  it  related  to  bismuth  chloride  ? 

SODIUM. 

Experiment  108.  —  Properties  of  Sodium.  Materials:  Sodium, 
pneumatic  trough  filled  with  water  as  usual,  litmus  paper,  filter  paper, 
tea  lead. 

Precaution.     Observe  the  precautions  as  in  Experiment  24. 

(a)  Examine  a  small  piece  of  sodium,  and  record  its  most  obvious 
physical  properties,  e.g.  color,  luster,  whether  hard  or  soft,  etc. 

(£)    Repeat  Experiment  24  (except  (*)). 

ANSWER  : 

(1)  Is  sodium  heavier  or  lighter  than  water? 

(2)  What  properties  show  that  it  is  a  metal? 


Sodium.  523 

(3)  Is  it  harder  or  softer  than  most  metals  ? 

(4)  What  is  the  test  for  sodium  ? 

Experiment  109.  —  Preparation  and  Properties  of  Sodium  Hy- 
droxide. Materials :  Sodium  carbonate,  lime,  zinc  sulphate  solution, 
iron  (or  tin)  dish,  file. 

Dissolve  25  grams  of  sodium  carbonate  in  150  cubic  centimeters  of 
water  and  heat  gently  in  an  iron  dish  (an  ordinary  iron  spider  is  well 
adapted  for  this  work).  Meanwhile  slake  10  grams  of  lime  by  adding 
just  enough  water  to  make  a  milky  liquid  —  "  milk  of  lime."  Add  the 
milk  of  lime  to  the  sodium  carbonate  solution  and  boil  for  several  min- 
utes, stirring  constantly  with  a  file.  Let  the  precipitate  settle,  remove 
a  little  liquid  with  a  small  tube,  and  if  it  effervesces  with  hydrochloric 
acid,  add  more  milk  of  lime  and  boil;  if  not,  pour  the  liquid  into  a  con- 
venient vessel,  let  it  stand  for  a  few  minutes  or  until  the  solid  settles ; 
then  pour  the  liquid  down  a  glass  rod  (  see  Int.  §  6  (i)  )  into  a  bottle. 
This  solution  of  sodium  hydroxide  may  be  evaporated  to  dryness,  and 
the  solid  product  tested  and  the  remainder  preserved,  or  the  solution 
may  be  tested  at  once  as  follows :  — 

(a)    Rub  a  little  between  the  fingers  and  describe  the  feeling. 

($)    Apply  the  litmus  test.     Is  it  acid  or  alkaline?     Is  it  markedly  so? 

(c)  Add  a  little  to  a  zinc  sulphate  solution,  and  shake.  The  precipi- 
tate is  zinc  hydroxide.  Describe  it.  Now  add  an  excess  of  sodium 
hydroxide,  and  shake.  Describe  the  result.  The  excess  of  sodium 
hydroxide  forms  soluble  sodium  zincate.  This  behavior  of  zinc  com- 
pounds is  the  test  for  an  hydroxide. 

(d}  How  do  sodium  compounds  affect  a  colorless  flame?  Try  it,  if 
in  doubt. 

Exercises  for  Review. 

1.  How  does  sodium  carbonate  act  when  exposed  to  air?     Sodium 
sulphate?     Sodium  hydroxide? 

2.  How  does  sodium  carbonate  solution  affect  litmus  paper?     How 
does  its  action  compare  with  that  of  other  salts,  sodium  chloride  for 
example  ? 

3.  What  marked  property  has  sodium  carbonate,  and  how  is  this 
property  utilized  ? 

4.  How  does  sodium  bicarbonate  interact  with  tartaric  acid  ?     Would 
the  action  be  the  same  with  other  acids? 


Experiments. 

5.  "  Sodium  thiosulphate  forms  a  supersaturated  solution."     Explain 
and  illustrate  this  statement. 

6.  How  does  sodium  chloride  interact  with  sulphuric  acid  ?     Sodium 
nitrate  with  sulphuric  acid  ? 

POTASSIUM. 

Experiment  110. — Properties  of  Potassium.     Materials:  Potas- 
sium, pneumatic  trough  filled  as  usual,  litmus  paper. 
Precaution.     Observe  the  same  precaution  as  in  using  sodium. 

(a)  Examine  a  very  small  piece  of  freshly  cut  potassium,  and  record 
its  most  obvious  physical  properties.     Touch  it  slightly.     Does  it  sug- 
gest caustic  potash  and  soda? 

(b)  Drop  a  small  piece  of  potassium  on  the  water  in  a  pneumatic 
trough.     Stand  just   near  enough   to   see   the   action.     Describe   the 
action.     How  does  it  differ  from. the  action  of  sodium?     Test  the  water 
as  in  Experiment  24  (d). 

From  what  has  already  been  learned  about  sodium  and  potassium, 
predict  the  main  chemical  change  observed  in  (£). 

ANSWER  : 

(1)  Is  potassium  heavier  or  lighter  than  water? 

(2)  What  properties  suggest  that  it  is  a  metal? 

(3)  How  does  it  resemble  and  differ  from  sodium? 

(4)  How  does  potassium  color  a  flame  ? 

(5)  What  is  the  test  for  potassium? 

Experiment  111.  —  Preparation  and  Properties  of  Potassium  Hy- 
droxide. Materials:  Potassium  carbonate,  lime,  zinc  sulphate  solu- 
tion, iron  (or  tin)  dish,  file. 

Proceed  as  in  Experiment  109,  but  use  potassium  carbonate  instead  of 
sodium  carbonate.  Test  as  in  the  case  of  sodium  hydroxide. 

Experiment  112. —Preparation  and  Properties  of  Potassium 
Carbonate.  Materials:  Cream  of  tartar,  wood  ashes,  litmus  paper, 
hydrochloric  acid,  iron  sand  bath  pan,  mortar  and  pestle. 

(a)  Heat  strongly  5  grams  of  cream  of  tartar  —  acid  potassium  tar- 
trate  —  in  an  iron  pan  in  the  hood  until  the  residue  is  nearly  white. 
Grind  this  solid  with  water  in  a  mortar,  and  filter.  Test  the  filtrate 
(i)  with  both  kinds  of  litmus  paper,  (2)  for  potassium,  and  (3)  for  a 
carbonate.  Record  the  results- 


Copper.  525 

(£)  Fill  a  test  tube  half  full  of  wood  ashes,  add  half  the  volume  of 
water,  shake,  and  warm  gently.  Filter,  and  test  the  filtrate  as  in  (#). 
If  test  (3)  is  not  decisive,  repeat  the  experiment  on  a  larger  scale. 
Record  the  results. 

(c)  Expose  a  little  potassium  carbonate  to  the  air  for  an  hour  or  more. 
Describe  the  result.  How  does  its  behavior  compare  with  that  of  sodium 
carbonate  under  the  same  conditions  ? 

ANSWER  : 

(1)  What  is  the  source  of  cream  of  tartar? 

(2)  What  do  (a)  and  (£)  show  about  the  distribution  of  potassium? 
Of  its  assimilation  by  plants? 

(3)  What  is  the  literal  meaning  of  the  word  potash  ? 


Exercises  for  Review. 

1.  What  does  potassium  chlorate  yield  when  heated? 

2.  Does  potassium  chlorate  dissolve  readily  in  cold  water?     In  hot 
water  ? 

3.  What  is  formed  by  heating  potassium  bromide  with  manganese 
dioxide  and  sulphuric  acid  ? 

4.  Apply  question  3  to  potassium  iodide. 

5.  What  happens  to  potassium  hydroxide  when  exposed  to  air  ?     To 
potassium  carbonate? 

6.  Of  what  important  mixture  is  potassium  nitrate  an  ingredient  ? 


COPPER. 

Experiment  113.  —  Physical  Properties  of  Copper. 

Examine  several  forms  of  copper  —  wire,  sheet,  filings,  etc.  —  and 
state  the  most  obvious  physical  properties. 

ANSWER  : 

% 

(1)  Is  copper  a  good  conductor  of  heat  ?     Of  electricity  ?     On  what 
evidence  is  your  answer  based  ? 

(2)  Is  copper  ductile  ?     Malleable  ?     Brittle  ?     Tough  ?     Hard  or 
soft? 

(3)  What  happens  to  copper  when  heated  ?    When  exposed  to  the 
air  ? 


526  Experiments. 

Experiment  114.  —  Tests  for  Copper.  Materials :  Copper  wire, 
copper  sulphate  solution,  ammonium  hydroxide,  acetic  acid,  potassium 
ferrocyanide  solution. 

(/i)  Heat  a  copper  wire  in  the  Bunsen  flame.  The  color  is  charac- 
teristic of  copper  and  its  compounds,  though  not  a  conclusive  test, 
since  the  same  color  is  produced  by  other  substances. 

(^)  Add  a  few  drops  of  ammonium  hydroxide  to  copper  sulphate 
solution,  and  observe  the  result;  now  add  an  excess  of  ammonium 
hydroxide.  The  blue  solution  is  a  characteristic  and  decisive  test 
for  copper. 

(c)  Add  to  a  test  tube  one  fourth  full  of  water  an  equal  volume  of 
copper  sulphate  solution,  and  shake ;  then  add  a  few  drops  of  acetic 
acid  and  of  potassium  ferrocyanide  solution.  The  brown  precipitate 
is  copper  ferrocyanide. 

Experiment  115.  —  Interaction  of  Copper  with  Metals.  Mate- 
rials: Copper  wire,  iron  nail,  zinc,  solutions  of  copper  sulphate  and 
any  mercury  compound. 

(a)  Put  a  clean  copper  wire  into  a  solution  of  any  mercury  compound. 
After  a  short  time,  remove  the  wire  and  wipe  it  with  a  soft  cloth  or 
paper.  Describe  the  change.  What  has  become  of  some  of  the  copper? 

(£)  Put  in  separate  test  tubes  half  full  of  copper  sulphate  solution 
a  bright  iron  nail  and  a  strip  of  clean  zinc.  After  a  short  time  remove 
the  metals  and  examine  them.  What  is  the  deposit  ?  What  has  become 
of  some  of  the  zinc  and  iron?  Does  the  final  color  of  the  solution 
indicate  any  chemical  change  ?  How  would  you  prove  the  answer  to 
the  last  question  ? 

Exercises  for  Review. 

1.  What  happens  to  a  crystal  of  copper  sulphate  when  heated  ? 

2.  You  are  given  a  blue  solution  supposed  to  be  copper  sulphate. 
State  how  you  would  prove  it. 

3.  What  are  formed  by  the  interaction  of  copper  and  nitric  acid? 

4.  What  color  have  many  copper  compounfls? 

5.  "  Brass  is  an  alloy  of  copper."     State  how  you  would  prove  this. 

SILVER  AND  GOLD. 

Experiment  116.  —  Preparation  of  Silver.  Materials :  (#)  Sil- 
ver nitrate  solution,  mercury,  evaporating  dish ;  (b)  ten-cent  piece. 


Silver  and  Gold.  527 

concentrated  nitric  acid,  hydrochloric  acid,  sulphuric  acid,  zinc,  evapo- 
rating dish,  charcoal,  sodium  carbonate,  blowpipe. 

Prepare  silver  by  one  or  both  of  the  following  methods :  — 

(a}  Fill  a  porcelain  dish  half  full  of  silver  nitrate  solution,  and  add  a 
few  drops  of  mercury.  Allow  the  whole  to  stand  undisturbed  for  a  day 
or  more,  and  then  examine.  The  delicate  crystals  attached  to  the 
mercury  are  silver.  Pick  them  out  with  the  forceps,  wash  well  with 
water,  and  preserve  them  for  Experiment  117. 

(£)  Dissolve  a  ten-cent  piece  in  10  cubic  centimeters  of  concentrated 
nitric  acid,  dilute  with  an  equal  volume  of  water,  and  add  hydrochloric 
acid  until  the  precipitation  is  complete.  Let  the  precipitate  settle. 
Then  filter,  and  wash  until  the  filtrate  is  neutral.  If  convenient,  let 
the  precipitate  dry ;  if  not,  scrape  half  from  the  opened  paper  with  a 
knife,  put  it  in  a  porcelain  dish,  cover  with  dilute  sulphuric  acid,  and 
add  a  piece  of  zinc ;  put  the  other  half  in  a  cavity  at  the  end  of  a  piece 
of  charcoal,  cover  with  sodium  carbonate,  and  reduce  it  with  a  blowpipe 
flame.  In  the  first  case,  the  silver  will  collect  as  a  grayish  powder; 
remove  any  excess  of  zinc,  filter,  wash  with  water  and  dry  the  residue. 
It  may  be  preserved  as  a  powder,  or  fused  into  a  bead  with  a  blowpipe 
flame.  In  the  second  case,  minute  globules  of  silver  will  appear  on  the 
charcoal ;  scrape  them  together  and  fuse  into  a  single  bead.  Preserve  it. 

Experiment  117. — Properties  of  Silver. 

Examine  the  silver  formed  in  Experiment  116,  and  state  briefly  its 
most  obvious  properties. 

Experiment  118.  —  Test  for  Silver. 

Devise  a  test  for  combined  silver,  based  upon  pYevious  experiments. 
Verify  it. 

«• 
Exercises  for  Review. 

1.  What  is  formed  by  the  interaction  of  silver  nitrate  and  potassium 
chloride?     Potassium   bromide?     Potassium   iodide?      How  do   the 
products  differ  ? 

2.  What  caused  the  blue  filtrate  in  Experiment  116  (b)  ?     What 
metal  besides  silver  does  a  ten-cent  piece  contain  ? 

3.  What  compound  is  formed  when  silver  tarnishes? 

4.  What  is  the  test  for  a  chloride  ? 


528  Experiments. 

Experiment  119.  —  Interaction  of  Gold  and  Aqua  Regia  and 
the  Test  for  Gold.  Materials :  Gold  leaf,  concentrated  nitric  and 
hydrochloric  acids,  stannous  chloride  solution. 

Prepare  a  solution  of  gold  chloride  according  to  Experiment  58,  using 
as  small  a  volume  of  the  acids  as  possible.  Dilute  with  water,  and  then 
slowly  add  a  dilute  solution  of  stannous  chloride.  A  precipitate  is 
produced,  varying  in  color  from  faint  purple  to  black  according  to  the 
conditions.  This  precipitate  is  supposed  to  be  finely  divided  gold,  and 
is  called  Purple  of  Cassius ;  its  formation  is  the  test  for  gold. 

CALCIUM. 

Experiment  120.  —  Tests  for  Calcium. 

(«)  Subject  calcium  chloride  to  the  flame  test.     Record  the  result. 
(b)  Repeat  Experiment  12  (d). 

Experiment  121.  —  The  ' '  Setting  ' '  of  Plaster  of  Paris.  Materials: 
Plaster  of  Paris,  block  of  wood. 

Mix  a  little  plaster  of  Paris  with  enough  water  on  a  block  of  wood  to 
form  a  thin  paste.  Let  it  stand  undisturbed  for  a  few  minutes,  and 
then  examine.  Describe  the  change.  How  is  this  property  utilized  ? 

Exercises  for  Review. 

1.  Describe   calcium   chloride.     How  does  it  act  when  exposed  to 
the  air  ?     How  would  you  show  that  it  (a)  contains  calcium  and  (6)  is 
a  chloride  ? 

2.  Describe  lime.     What  effect  does  water  have  upon  it  ? 

3.  What  compounds  are  produced  by  the  interaction  of  calcium  car- 
bonate and  hydrochloric  acid  ? 

4.  What  is  limewater  ?     Milk  of  lime  ? 

5.  What  is  formed  when  carbon  dioxide  is  passed  into  limewater  ? 
Into  sodium  hydroxide  ? 

6.  What  compounds  are  formed  by  heating  calcium  carbonate  ? 

7.  What  happens  when  an  excess  of  carbon  dioxide  is  passed  into 
limewater  ? 

8.  What  is  acid  calcium  carbonate  ?     How  can  it  be  changed  into 
normal  calcium  carbonate  ? 

9.  What  compounds  are  formed  by  the  interaction  of  calcium  fluoride 
and  sulphuric  acid  ? 


Barium.  529 

10.  What  is  calcium  hypochlorite  ?     For  what  is  it  used  ? 

1 1 .  What  happens  to  crystallized  calcium  sulphate  (selenite)  when 
heated  ? 

12.  Calcium  sulphate  is  nearly  insoluble  in  water;    how  can  it  be 
shown  to  be  a  sulphate  ? 

13.  What  is  tricalcium  phosphate  ?     Superphosphate  of  lime  ? 

14.  What  is  the  scientific  name  of  lime,  limestone,  marble,  gypsum, 
fluor  spar,  limewater,  slaked  lime,  quicklime,  bleaching  powder,  chloride 
of  lime  ? 

STRONTIUM. 
Experiment  122.  —  Test  for  Strontium. 

Dip  a  platinum  test  wire  (or  a  glass  rod)  into  a  solution  of  strontium 
nitrate,  and  hold  it  in  the  Bunsen  flame.  Describe  the  result,  after 
several  trials. 

Experiment  123.  —  Red  Fire.  Materials:  Strontium  nitrate,  pow- 
dered potassium  chlorate,  powdered  shellac,  iron  pan  or  brick. 

Mix  carefully  small  and  equal  (in  bulk)  quantities  of  the  three  sub- 
stances on  a  sheet  of  paper.  Place  the  mixture  on  a  sand-bath  pan  or 
a  brick  in  the  hood,  and  light  it  with  a  Bunsen  burner.  Describe  the 
result. 

BARIUM. 

(Compounds  of  Barium  are  Poisonous?) 
Experiment  124.  —  Tests  for  Barium. 

(a)  Repeat  Experiment  122,  using  a  solution  of  barium  chloride  or 
barium  nitrate.      Be  sure  the  test  wire  (or  rod)  is  clean.     Describe  the 
result. 

(b)  Devise  a  test.     (Suggestion.     What  is  the  test  for  a  sulphate  ?) 

Experiment  125.  —  Green  Fire. 

Repeat  Experiment  123,  using  barium  nitrate  instead  of  strontium 
nitrate. 

Exercises  for  Review. 

1.  State  the  properties  of  barium  sulphate. 

2.  How  does  barium  chloride  behave  toward  litmus  ? 

3.  What  industrial  use  have  the  barium  oxides  ? 


530  Experiments. 


MAGNESIUM. 

Experiment  126.  —  Properties  of  Magnesium. 
Examine  a  piece  of  magnesium  and  state  briefly  its  most  obvious 
properties. 

Experiment  127.  —  Tests  for  Magnesium.  Materials:  Solutions 
of  magnesium  sulphate  (or  chloride),  ammonium  chloride,  ammonium 
hydroxide,  disodium  phosphate,  and  cobaltous  nitrate ;  magnesium 
oxide,  charcoal,  blowpipe. 

(a)  To  a  solution  of  magnesium  sulphate  (or  chloride)  add  succes- 
sively solutions  of  ammonium  chloride,  ammonium  hydroxide,  and  di- 
sodium phosphate.     A  precipitate  of  ammonium  magnesium  phosphate 
is  formed.     It  is  voluminous  at  first,  but  finally  crystalline.     It  is  soluble 
in  acids.     Try  it. 

(b)  Put  a  little  powdered  magnesium  oxide  in  a  cavity  at  the  end 
of  a  piece  of  charcoal,  moisten  with  water,  and  heat  intensely  in  a 
blowpipe  flame.     Cool,  and  moisten  with  a  drop  of  cobaltous  nitrate 
solution.     Heat  again,  and  when  cool  observe  the  color.     If  the  experi- 
ment has  been  conducted  properly,  a  pink  or  pale  flesh-colored  residue 
coats  the  charcoal.     Describe  the  result. 

Exercises  for  Review. 

1.  What  compound  is  formed  by  burning  magnesium  in  air  or  in 
oxygen?     Describe  it. 

2.  How  was  magnesium  utilized  in  the  discovery  of  argon  ? 

ZINC. 
Experiment  128.  —  Properties  of  Zinc. 

Examine  a  piece  of  zinc  and  record  its  most  obvious  properties. 
ANSWER  : 

1.  What  happens  to  zinc  when  it  is  heated?     Describe  and  name 
the  product. 

2.  Is  zinc  hard  or  soft?     Malleable?     Ductile?    Brittle?     Tough? 
Does  it  melt  readily? 

Experiment  129.  —  Tests  for  Zinc.  Materials :  Zinc  oxide,  cobalt- 
ous nitrate  solution,  charcoal,  blowpipe. 


Mercury.  531 

(a)  Recall  or  devise  a  simple  test  for  combined  zinc.      (Suggestion. 
See  Exp.  109  (<:).) 

(b)  Recall  or  repeat  the  action  of  zinc  when  heated  in  the  oxidizing 
flame.     (See  Exp.  79  II  (a).) 

(<;)  Fill  a  small  cavity  at  one  end  of  a  piece  of  charcoal  with  zinc 
oxide,  moisten  with  water,  and  heat  strongly  in  the  blowpipe  flame. 
Cool,  and  moisten  with  a  drop  of  cobaltous  nitrate  solution,  then  heat 
again.  Cool  and  examine.  A  green  incrustation  is  caused  by  zinc 
compounds. 

Experiment  130.  —  Interaction  of  Zinc  and  Metals.  Materials  : 
Sheet  zinc,  solutions  of  copper  sulphate,  lead  nitrate,  mercurous  nitrate. 

(a)  Repeat  Experiment  115  (£). 

(b)  Repeat  (a)  using  lead  nitrate  solution. 

(c)  Repeat  (#)  using  the  mercury  salt  solution.     Examine  after  a 
short  time,  and  describe.     What  is  amalgamated  zinc,  and  for  what  is  it 
used? 

Exercises  for  Review. 

1.  What  are  formed  by  the  interaction  of  zinc  and  sulphuric  acid? 
Of  zinc  and  nitric  acid  ? 

2.  What   is   formed   by   the   interaction  of  a  zinc  salt  and  a  little 
sodium  hydroxide  solution  ?     An  excess  of  the  alkali  ? 

CADMIUM. 

Experiment  131.  —  Test  for  Cadmium. 

Add  hydrogen  sulphide  water  to  a  test  tube  half  full  of  cadmium 
chloride  solution.  The  precipitate  is  cadmium  sulphide.  Describe  it. 
Let  it  settle,  pour  off  most  of  the  liquid,  fill  the  test  tube  half  full  with 
dilute  sulphuric  acid,  and  warm.  Describe  the  result. 

MERCURY. 

{Mercury  and  its  Compounds  are  Poisonous.} 
Experiment  132.  —  Properties  of  Mercury. 

(a)  Pour  a  drop  or  two  of  mercury  into  an  evaporating  dish.     Ex- 
amine the  mercury,  and  state  its  characteristic  properties.     Agitate  the 
dish,  and  describe  the  result.     Why  is  mercury  called  u  quicksilver  "  ? 

(b)  Lift  carefully  a  bottle  of  mercury.     Estimate  the  specific  gravity. 
Verify  the  estimate  by  consulting  a  book. 


532  Experiments. 

Experiment  133.  —  Tests  for  Mercury. 

(a)  What  is  a  simple  test  for  free  mercury? 

(b)  Recall  or  devise  a  test  for  combined  mercury.     Verify  it.     (Sug- 
gestion.    See  Exp.  115.) 

Experiment  134. — Properties  of  Mercurous  and  Mercuric 
Compounds.  Materials:  Solutions  of  mercurous  and  mercuric 
nitrate  ;  hydrochloric  acid,  ammonium  hydroxide. 

(a)  Mercurous.  Add  a  few  drops  of  hydrochloric  acid  to  a  little 
mercurous  nitrate  solution.  The  white  precipitate  is  mercurous  chlo- 
ride. Note  its  insolubility  in  water  and  in  dilute  hydrochloric  acid. 
Add  a  few  drops  of  ammonium  hydroxide.  The  black  precipitate  is 
mercurous  ammonium  chloride.  Its  formation  is  a  delicate  test  for 
mercury  in  mercurous  compounds. 

(b}  Mercuric.  Add  a  few  drops  of  hydrochloric  acid  to  a  little 
mercuric  nitrate  solution.  Compare  the  result  with  that  in  (a).  Add 
a  few  drops  of  ammonium  hydroxide,  or  enough  to  produce  a  decided 
change.  Compare  with  (a).  The  precipitate  is  mercuric  ammonium 
chloride. 

Exercises  for  Review. 

1.  Describe  the  effect  of  heat  on  red  oxide  of  mercury.     What  his- 
torical interest  has  this  experiment  ? 

2.  What  practical  use  has  mercury  ? 

3.  What  are  amalgams  ? 

4.  What  action  has  mercury  upon  gold  ? 

ALUMINIUM. 

Experiment  135.  —  Properties  of  Aluminium. 

(a)  Examine  a  piece  of  aluminium,  and  observe  its  properties.  Has 
it  any  "spring"  like  brass  ?  Is  it  ductile,  malleable,  soft,  hard,  tough, 
brittle  ?  Will  it  melt  in  the  Bunsen  flame  ?  Try  it. 

(b}  Compare  roughly  the  weight  of  a  piece  of  sheet  aluminium  with 
a  piece  of  pasteboard  or  glass  having  approximately  the  same  volume. 
Estimate  the  specific  gravity.  Verify  it  by  consulting  a  book. 

Experiment  136.— Action  of  Aluminium  with  Acids  and  Alka- 
lies. Materials:  Aluminium,  sulphuric  acid,  hydrochloric  acid,  sodium 
hydroxide  solution. 


Aluminium. 


533 


(«)  Add  a  small  piece  of  aluminium  to  separate  test  tubes  containing 
dilute  sulphuric  acid  and  concentrated  hydrochloric  acid.  Warm,  if 
necessary.  Describe  the  action.  Test  the  gas  evolved.  What  com- 
pound is  formed  in  each  case  ? 

(b}  Add  a  small  piece  of  aluminium  to  a  test  tube  half  full  of  dilute 
sodium  hydroxide  solution,  and  boil.  Test  any  gas  evolved.  If  only 
a  little  gas  is  liberated,  attach  a  simple  delivery  tube  and  collect  the 
gas  over  water. 

Other  acids  and  alkalies  act  similarly.  Draw  a  general  conclusion 
from  this  experiment. 

Experiment  137.  —  Preparation  and  Properties  of  Aluminium 
Hydroxide.  Materials:  Solutions  of  alum,  sodium  hydroxide,  am- 
monium sulphide,  and  cochineal ;  hydrochloric  acid  and  ammonium 
hydroxide. 

(a)  Add  slowly  a  little  sodium  hydroxide  solution  to  a  test  tube 
half  full  of  alum  solution.  The  gelatinous  precipitate  is  aluminium 
hydroxide.  Now  add  an  excess  of  the  alkali  to  one  half,  and  dilute 
hydrochloric  acid  to  the  other.  Describe  the  results. 

(b}  Add  a  little  solution  of  ammonium  sulphide  to  a  solution  of  alum. 
Describe  the  result.  The  precipitate  is  not  a  sulphide,  but  aluminium 
hydroxide,  because  aluminium  forms  no  sulphide  in  the  wet  way. 

(c}  Add  a  little  alum  solution  to  a  dilute  solution  of  cochineal,  then 
add  ammonium  hydroxide.  The  colored  product  is  called  carmine  lake. 
It  belongs  to  a  class  of  dyes  formed  by  the  combination  of  a  vegetable 
dye  and  a  metallic  hydroxide,  often  aluminium  hydroxide. 

Experiment  138.  —  Tests  for  Aluminium.  Materials  for  (c)  : 
Aluminium  sulphate ;  cobaltous  nitrate  solution,  blowpipe,  charcoal. 

(a}  What  is  a  simple  test  for  metallic  aluminium  ? 

(&)  Recall  or  devise  a  test  for  combined  aluminium.  Verify  it. 
How  can  aluminium  compounds  be  distinguished  from  those  of  zinc? 

(c)  Heat  a  little  aluminium  sulphate  on  charcoal  in  the  blowpipe 
flame.  Cool,  and  moisten  with  a  drop  of  cobaltous  nitrate  solution. 
Heat  again,  and  if  the  operation  has  been  conducted  properly,  a  blue 
residue  will  coat  the  charcoal.  This  color  is  characteristic  of  aluminium 
compounds.  Compare  this  result  with  the  action  of  other  metallic 
compounds  under  similar  circumstances. 


534  Experiments. 

Experiment  139.  —  Preparation  and  Properties  of  Common 
Alum.  Materials  for  (a) :  Aluminium  sulphate,  potassium  sulphate, 
evaporating  dish. 

(a)  Dissolve  about  10  grams  of  aluminium  sulphate  in  the  least  pos- 
sible amount  of  hot  water.  Dissolve  3  grams  of  potassium  sulphate  in 
the  same  way.  Mix  the  clear,  hot,  saturated  solutions  in  an  evaporat- 
ing dish,  and  allow  the  solution  to  cool  undisturbed.  Crystals  of 
potassium  alum  will  be  deposited.  Remove  the  best  ones;  dry  and 
examine.  Describe  them,  giving  color,  luster,  size,  and  crystal  form. 

(£)  Prove  by  actual  tests  that  (i)  they  are  a  sulphate,  and  (2)  they 
contain  aluminium  and  water  of  crystallization. 


TIN. 

Experiment  140.  —  Properties  of  Tin.  Examine  a  piece  of  tin,  and 
state  its  most  obvious  properties.  Is  it  malleable,  soft,  hard,  tough, 
brittle?  Will  it  melt  in  the  Bunsen  flame?  Try  it. 

Experiment  141. — Action  of  Tin  with  Acids.  Materials:  Tin, 
concentrated  nitric  and  hydrochloric  acids. 

(a)  Put  a  small  piece  of  tin  in  a  test  tube,  cover  with  concentrated 
hydrochloric  acid,  add  a  little  water,  and  heat  —  in  the  hood.  Heat 
gently  at  first,  and  when  action  begins  regulate  the  heat  accordingly. 
Most  of  the  tin  disappears,  soluble  stannous  chloride  being  formed. 
Save  this  solution  for  Experiment  142. 

(<£)  Treat  a  small  piece  of  tin  with  concentrated  nitric  acid  —  in  the 
hood.  It  is  advisable  to  stand  the  test  tube  in  the  rack  or  in  a  bottle  as 
soon  as  the  action  begins.  The  white,  amorphous  product  is  metastan- 
nic  acid.  How  does  the  action  of  nitric  acid  on  tin  differ  from  and 
resemble  its  action  on  other  metals,  zinc,  for  example? 

Experiment  142.  —  Tests  for  Tin.  Materials  for  (c)  :  Solutions  of 
mercuric  chloride  and  stannous  chloride. 

(a)  What  is  a  simple  test  for  metallic  tin? 

(£)  Recall  or  repeat  the  action  of  tin  when  heated  in  a  blowpipe 
flame. 

(c)  Add  a  few  drops  of  mercuric  chloride  solution  (poison)  to  a  little 
of  the  stannous  chloride  solution  prepared  in  Experiment  141 .  The  white 
precipitate  is  mercurous  chloride.  Add  a  little  more  stannous  chloride 


Lead.  535 

solution  and  heat  gently.     The  mercurous  chloride  is  reduced  finally  to 
mercury,  which  appears  as  a  grayish  powder. 

Experiment  143. — Deposition  of  Tin. 

Put  a  strip  of  zinc  in  a  slightly  acid  solution  of  stannous  chloride. 
Examine  after  a  short  time.  The  tin  will  be  found  adhering  to  the 
zinc  as  a  grayish  black  deposit;  sometimes  bright  scales  are  also  seen. 
What  has  become  of  the  zinc? 

LEAD. 

Experiment  144.  —  Properties  of  Lead. 

(a}  Examine  a  piece  of  freshly  cut  lead  and  state  its  most  obvious 
physical  properties. 

(£)  Estimate  its  specific  gravity.  Verify  your  estimate  by  consulting 
a  book. 

(Y)  Draw  a  piece  of  lead  across  a  sheet  of  white  paper,  and  describe 
the  result. 

(d}  Is  lead  easily  melted?    Try  it. 

ANSWER  : 

(1)  What  happens  to  lead  when  exposed  to  the  air? 

(2)  What  properties  adapt  lead  for  its  extensive  use? 

(3)  What  is  "black  lead"? 

(4)  Is  there  lead  in  a  lead  pencil? 

Experiment  145. — Tests  for  Lead.  Materials  for  (Y),  (*/),  (*)  : 
Lead  nitrate  and  potassium  dichromate  solutions,  sulphuric  acid,  hydro- 
chloric acid. 

(a)  Recall  or  repeat  the  reduction  of  lead  oxide  in  the  blowpipe 
flame.  (See  Exp.  79.) 

(d)  Recall  or  repeat  the  action  of  hydrogen  sulphide  with  the  solu- 
tion of  a  lead  compound.  (See  Exp.  95  (£)•) 

(£•)  Add  dilute  hydrochloric  acid  to  a  little  lead  nitrate  solution  until 
precipitation  ceases.  Note  the  insolubility  of  the  lead  chloride  which  is 
formed.  Warm  gently  as  long  as  any  decided  change  occurs.  Describe 
the  action.  This  is  characteristic  of  lead  chloride  and  permits  its 
separation  from  the  chlorides  of  silver  and  of  mercury  (in  the  -ous 
condition). 


53  6  Experiments. 

(d)  Add  dilute  sulphuric  acid  to  a  little  lead  nitrate  solution  until 
precipitation  ceases.     The   precipitate  is  lead  sulphate.     Observe  its 
properties.     Is  it  soluble  in  hot  water  ?    Try  it. 

(e)  Repeat  (</),  using  potassium  dichromate  solution  instead  of  sul- 
phuric acid.    The  precipitate  is  lead  chromate.     Describe  it,  especially 
the  color. 

Experiment  146.  —  Deposition  of  Lead. 

Repeat  Experiment  130  (b). 

Experiment  147.  —  Properties  of  Lead  Oxides.  Materials:  Lead 
monoxide,  dioxide,  and  tetroxide,  nitric  acid. 

(a)  Examine  the  three  oxides  and  tabulate  their  most  obvious  physi- 
cal properties,  stating  the  exact  chemical  name  and  formula  and  the 
popular  name  of  each  oxide. 

(b}  Recall  the  experiment  in  which  lead  was  heated  in  the  oxidizing 
flame,  especially  the  color  of  the  coating.  What  oxide  of  lead  is 
thereby  formed  ? 

(c)  Warm  a  little  lead  tetroxide  with  dilute  nitric  acid.  The  solid 
product  is  lead  dioxide.  Describe  it. 

EXERCISES  : 

(1)  How  might  lead  tetroxide  be  prepared  ? 

(2)  If  lead  tetroxide  is  heated  strongly,  lead  monoxide  is  formed. 
What  does  this  fact  reveal  about  the  stability  of  lead  tetroxide  ? 

(3)  When  lead  dioxide  and  concentrated  hydrochloric  acid  are  mixed 
and  heated,  chlorine  is  evolved.     Complete  the  equation  — 

PbO2  +  4  HC1  =  PbCl2  +  2  H2O  + 

How  does  this  interaction  resemble  that  of  manganese  dioxide  and 
hydrochloric  acid  ? 

Experiment    148.  —  Properties    of    Certain    Lead    Compounds. 

Materials:    Lead   nitrate,  lead   carbonate,  galena,  taper,   mortar  and 
pestle. 

(#)  Put  a  crystal  of  lead  nitrate  in  a  test  tube  provided  with  a 
holder,  hold  in  a  horizontal  position,  and  heat  strongly  in  the  upper 
part  of  the  Bunsen  flame.  Describe  the  result.  What  is  the  most 
obvious  product  ?  After  most  of  this  product  has  passed  out  of  the 
test  tube,  thrust  well  into  the  test  tube  a  taper  with  a  spark  on  the 


Chromium.  537 

end.  Describe  the  result.  What  other  gas  is  present  ?  The  solid 
product  is  lead  monoxide.  Summarize  the  behavior  of  lead  nitrate 
when  heated. 

(b)  Examine  lead  carbonate,  and  state  its  most  obvious  properties. 
What  is  its  common  name  ?    Prove   that  it  is  a  carbonate  and  con- 
tains lead.     (Suggestion.     Treat  with  hydrochloric  acid.) 

(c)  Examine  a  lump  of  galena,  and  state  its  most  obvious  proper- 
ties.   Pulverize  it  in  a  mortar.    What  additional  property  is  revealed  ? 
Prove  that  it  is  a  sulphide  and  contains  lead.    (Suggestion.    See  Exps. 
93  and  145.) 

CHROMIUM. 

Experiment  149.  —  Tests  for  Chromium.  Materials:  Borax, 
chrome  alum,  potassium  carbonate,  potassium  nitrate,  acetic  acid,  nitric 
acid,  sodium  hydroxide  solution,  lead  nitrate  solution,  potassium  dichro- 
mate  solution,  platinum  test  wire,  piece  of  porcelain,  forceps. 

(a}  Prepare  a  borax  bead  (see  Exp.  102),  touch  it  with  a  minute 
quantity  of  chrome  alum,  and  heat  in  both  the  oxidizing  and  reducing 
flame.  Describe  the  result. 

(#)  Mix  equal  small  quantities  of  potassium  carbonate,  potassium 
nitrate,  and  powdered  chrome  alum,  place  the  mixture  on  a  piece  of 
porcelain,  and  hold  it  with  the  forceps  in  the  upper  Bunsen  flame  so 
that  the  mixture  will  fuse.  A  yellow  mass,  due  to  the  presence 
of  potassium  chromate,  results.  If  the  color  is  not  decided,  dissolve 
the  mass  in  water,  add  acetic  acid,  slowly  at  first,  and  boil  to  expel  the 
carbon  dioxide.  Add  a  few  drops  of  lead  nitrate  solution  to  a  portion, 
and  yellow  lead  chromate  is  precipitated.  If  the  precipitate  is  white, 
it  is  lead  carbonate,  and  shows  that  not  all  the  potassium  carbonate 
was  decomposed,  as  intended. 

(c)  Add  lead  nitrate  solution  to  potassium  dichromate  solution. 
Name  and  describe  the  precipitate.  Try  the  solubility  of  the  precipitate 
in  acetic  acid,  dilute  nitric  acid,  and  sodium  hydroxide.  Describe  the 
result. 

Experiment  150.  —  Properties  of  Chromates.  Materials :  Potas- 
sium chromate  and  dichromate,  concentrated  hydrochloric  acid, 
potassium  hydroxide  solution. 

(#)  Examine  crystals  of  potassium  chromate  and  dichromate,  and 
state  their  characteristic  properties.  Make  a  dilute  solution  of  each, 
and  compare  the  colors.  Save  for  (c)  and  (*/). 


53  8  Experiments. 

(b)    State  the  properties  of  lead  chromate. 

(V)  Add  a  few  drops  of  concentrated  hydrochloric  acid  to  the  solution 
of  potassium  chromate  prepared  in  (a),  and  observe  the  change  in 
color.  Describe  it.  Compare  with  the  color  of  the  potassium  dichro- 
mate  solution  prepared  in  (a).  Draw  a  conclusion. 

(d)  Add  potassium  hydroxide  solution  to  the  solution  of  potassium 
dichromate  prepared  in  (a)  until  a  change  of  color  is  produced.      De- 
scribe  the   color.     Compare   with   the   potassium   chromate  solution. 
Draw  a  conclusion. 

(e)  The  chromates  are  oxidizing  agents.      Add  a  few  drops  of  con- 
centrated hydrochloric  acid  to  powdered  potassium  chromate  and  dichro- 
mate in  separate  test  tubes.      Chlorine  is  evolved.     Where  did  it  come 
from  ?     By  what  general  chemical  change  ? 


Experiment  151.  —  Reduction  of  Chromates  to  Chromic  Com- 
pounds. Materials:  Potassium  dichromate  solution,  concentrated 
hydrochloric  acid,  alcohol. 

Add  to  a  few  cubic  centimeters  of  potassium  dichromate  solution  a 
little  concentrated  hydrochloric  acid  and  a  few  drops  of  alcohol.  Warm 
gently.  Two  important  changes  occur.  The  chromate  is  reduced  to 
chromic  chloride  which  colors  the  solution  green ;  the  alcohol  is  oxi- 
dized to  aldehyde,  which  is  detected  by  its  peculiar  odor. 


Experiment  152.  —  Preparation  and  Properties  of  Chromic 
Hydroxide.  Materials:  Ammonium  sulphide,  solutions  of  sodium 
hydroxide  and  chrome  alum. 

(a)  Add  a  little  sodium  hydroxide  solution  to  a  solution  of  chrome 
alum.  The  precipitate  is  chromic  hydroxide.  Describe  it.  Add  an 
excess  of  sodium  hydroxide  solution,  and  shake.  Describe  the  result. 
Boil,  and  state  the  result. 

(d)  Add  a  little,  and  then  an  excess,  of  ammonium  sulphide  to  a 
solution  of  chrome  alum.  Compare  the  result  with  that  in  (a) .  Does 
chromium  form  a  sulphide  ? 

ANSWER  : 

How  can  chromic  hydroxide  be  distinguished  from  aluminium 
hydroxide  ? 


Manganese.  539 

Experiment  153.  —  Properties  of  Chrome  Alum. 

(«)  Examine  chrome  alum  and  state  its  most  obvious  physical 
properties. 

(£)  Prove  that  chrome  alum  is  a  sulphate,  and  that  it  contains 
chromium  and  water  of  crystallization. 


MANGANESE. 

Experiment  154.  —  Tests  for  Manganese.  Materials  for  (b)  and 
(c) :  Manganese  dioxide,  potassium  carbonate,  potassium  nitrate,  am- 
monium sulphide,  manganese  sulphate  solution,  hydrochloric  acid, 
acetic  acid,  ammonium  hydroxide. 

(a)  Subject  a  minute  quantity  of  manganese  dioxide  to  the  borax 
bead  test,  and  note  the  color  of  the  bead  after  heating  in  each  flame. 
(See  Exp.  102.) 

(b)  Fuse,   on   a    piece    of   porcelain,    a    little    manganese    dioxide 
mixed  with  potassium  carbonate  and  potassium  nitrate.     The  green 
mass  is  a  test  for  manganese.     It  is  due  to  the  presence  of  potassium 
manganate. 

(c)  Add  ammonium  sulphide  to  manganese  sulphate  solution.     The 
flesh-colored  precipitate  is  manganese  sulphide.     Compare  with  other 
sulphides  as  to  color  (see  Exp.  95).    Divide  it  into  two  parts.     Add 
hydrochloric  acid  to  one,  and  acetic  acid  to  the  other,  then  add  an 
excess  of  ammonium  hydroxide  to  each.     Draw  a  conclusion  regarding 
the  solubility  of  manganese  sulphide. 

Experiment  155.  —  Oxidation  with  Potassium  Permanganate. 

Materials:  Potassium  permanganate,  sulphuric  acid,  ferrous  sulphate, 
filter  paper. 

(a)  Add  a  few  drops  of  sulphuric  acid  to  a  weak  solution  of  fresh 
ferrous  sulphate ;  then  add,  drop  by  drop,  a  dilute  solution  of  potas- 
sium permanganate.     Its  color  is  changed,  owing  to  the  loss  of  oxygen  ; 
the  latter  converts  the  ferrous   to   ferric  sulphate.      The   decomposi- 
tion of  the  permanganate  also  causes  the  formation  of  potassium  and 
manganese  sulphates. 

(b)  Pour  a  solution  of  potassium  permanganate  upon  a  piece  of  filter 
paper.     Describe  and  explain  the  result. 


540  Experiments. 

Exercises  for  Review. 

1.  Describe  manganese  dioxide.    Name  five  elements  in  whose  prepa- 
ration manganese  dioxide  is  used.     Is  manganese  dioxide  an  oxidizing 
agent  ? 

2.  Describe  potassium  permanganate.     What  can  be  said  of  its  solu- 
bility in  water  ?     In  what  previous  experiment  has  it  been  used  ? 

3.  What  is  the  formula  of  potassium   permanganate  ?     Does   the 
formula  suggest  its  oxidizing  power  ? 

IRON. 

Experiment  156.  —  Properties  of  Iron.  Materials:  Cast  and 
wrought  iron,  steel,  magnet,  iron  wire,  iron  powder. 

(a)  Examine  cast  iron,  wrought  iron,  and  steel,  and  state  their  most 
obvious  physical  properties 

(£)  Try  the  action  of  a  magnet  on  each.     Describe  the  result. 

(c)  Drop  a  pinch  of  iron  powder  into  the  Bunsen  flame.  Hold  a 
piece  of  fine  iron  wire  in  the  flame.  Describe  the  results,  and  draw  a 
conclusion. 

Experiment  157.  —  Properties  of  Ferrous  Compounds.  Mate- 
rials: Iron  powder  (or  filings),  hydrochloric  acid,  solutions  of  sodium 
hydroxide,  potassium  ferricyanide,  potassium  thiocyanate,  potassium 
ferrocyanide. 

Put  a  few  grams  of  iron  powder  in  a  test  tube,  add  about  10  cubic 
centimeters  of  dilute  hydrochloric  acid,  and  warm  gently ;  ferrous  chlo- 
ride is  formed  (in  solution).  Proceed  as  follows  :  (i)  Pour  a  little  into  a 
test  tube  one  third  full  of  sodium  hydroxide  solution.  The  precipitate  is 
ferrous  hydroxide.  Watch  the  changes  in  color.  To  what  are  the 
changes  due  ?  (2)  Add  a  second  portion  to  potassium  ferricyanide 
solution.  The  precipitate  is  ferrous  ferricyanide.  Describe  it.  (3)  Add 
a  third  portion  to  potassium  thiocyanate  solution.  If  ferric  salts  are 
absent,  no  change  results.  (4)  Add  a  fourth  portion  to  potassium 
ferrocyanide  solution.  The  precipitate  is  ferrous  ferrocyanide.  Describe 
it.  Tabulate  the  results  as  described  in  the  next  experiment. 

Experiment  158.  —  Properties  of  Ferric  Compounds.  Materials : 
Ferric  chloride  solution  and  the  solutions  used  in  Experiment  157. 


Iron.  541 

To  a  little  ferric  chloride  solution  add  (i)  sodium  hydroxide  solu- 
tion. The  precipitate  is  ferric  hydroxide.  Describe  it.  Add  to  ferric 
chloride  solution  {2)  a  little  solution  of  potassium  ferricyanide.  Com- 
pare the  negative  result  with  (2)  in  Experiment  157.  Add  as  above  (3) 
a  little  solution  of  potassium  thiocyanate.  The  rich  wine-red  coloration 
is  caused  by  the  soluble  ferric  thiocyanate.  This  test  distinguishes  ferric 
from  ferrous  compounds.  Add  as  above  (4)  a  little  solution  of  potas- 
sium ferrocyanide.  The  precipitate  is  ferric  ferrocyanide.  Describe  it. 
Tabulate  the  results  of  these  two  experiments,  showing  the  behavior 
of  ferrous  and  ferric  compounds  under  the  same  conditions. 

Experiment  159.  —  Reduction  of  Ferric  Compounds.  Mate- 
rials: Ferric  chloride  solution,  zinc,  hydrochloric  acid. 

Put  a  piece  of  zinc  in  ferric  chloride  solution  made  slightly  acid  by 
hydrochloric  acid.  The  nascent  hydrogen  reduces  the  ferric  to  ferrous 
chloride.  After  the  operation  has  proceeded  for  about  fifteen  minutes, 
test  a  portion  of  the  liquid  for  a  ferrous  and  a  ferric  compound  by  Ex- 
periments 157  (2)  and  158  (3).  If  the  tests  are  not  conclusive,  continue 
the  reduction  and  test  again.  Describe  the  result. 

Experiment  160.  —  Oxidation  of  Ferrous  Compounds.  Mate- 
rials: Ferrous  sulphate,  hydrochloric  acid,  potassium  chlorate,  nitric 
acid. 

(a)  To  a  solution  of  fresh  or  freshly  washed  ferrous  sulphate  add  a 
little  hydrochloric  acid,  warm  gently,  and  then  add  a  few  crystals  of 
potassium  chlorate.     After  heating  a  short  time,  test  portions  of  the 
liquid  for  a  ferric  and  a  ferrous  compound. 

(b)  Add  10  cubic  centimeters  of  concentrated  nitric  acid,  drop  by 
drop,  to  a  hot  solution  of  ferrous  sulphate  to  which  a  little  sulphuric 
acid  has  been  added,  and  boil.     Test  portions  of  the  liquid  for  a  ferric 
and  a  ferrous  compound  as  in  Experiment  159. 

(c}  Recall  a  third  illustration  of  the  oxidation  of  a  ferrous  to  a  ferric 
compound.  Describe  it  briefly. 

Experiment  161.  — Properties  of  Certain  Iron  Compounds.  Ma- 
terials :  Ferrous  sulphate,  hematite,  li'monite,  magnetite,  pyrite,  siderite. 

(a}  Examine  a  crystal  of  ferrous  sulphate,  and  state  its  most  obvious 
properties.  Heat  it  gently  in  a  test  tube  inclined  mouth  downward. 
Describe  the  result.  Test  this  crystal  for  ferrous  and  ferric  compounds 


Experiments. 

as  in  Experiment  159.  State  and  explain  the  result.  What  is  the 
common  name  of  ferrous  sulphate? 

(£)  Examine  hematite,  limonite,  and  magnetite,  and  state  their  prop- 
erties. Draw  the  first  two  across  a  rough  sheet  of  paper  or  a  piece  of 
ground  glass,  and  describe  the  "streak  "  made  by  each.  What  is  the 
formula  of  each  compound?  Prepare  a  hydrochloric  acid  solution  of 
each  and  test  for  iron.  State  the  result. 

(V).  Examine  pyrite,  and  state  its  properties.  It  is  iron  disulphide. 
What  is  its  formula?  For  what  is  it  used?  For  what  is  it  sometimes 
mistaken  ? 

(d)  Examine  siderite,  and  state  its  properties.  It  is  ferrous  carbon- 
ate. What  is  its  formula?  Test  a  powdered  specimen  for  a  carbonate 
and  for  iron.  State  the  result. 

Exercises  for  Review. 

1.  What  happens  when  iron  is  (a)  treated  with  acids,  and  (b)  heated 
with  sulphur? 

2.  Describe  ferrous  sulphide.     What  are  formed  by  its  interaction 
with  warm  hydrochloric  acid? 

3.  What  happens   to   iron  when   it   is   placed   in  copper   sulphate 
solution  ? 

NICKEL  AND   COBALT. 

Experiment  162.  —  Test  for  Nickel. 

To  a  solution  of  nickel  chloride  add  sodium  hydroxide  to  alkaline 
reaction.  The  precipitate  is  nickelous  hydroxide.  Describe  it. 

Experiment  163.  — Test  for  Cobalt.     Repeat  Experiment  102  (a). 

ORGANIC  COMPOUNDS. 

Experiment  164.  —  Composition  of  Organic  Compounds,  Ma- 
terials: Turpentine,  alcohol,  camphor,  kerosene,  sugar,  starch,  flour, 
wood,  paper,  hair,  candle,  taper,  gelatine,  mustard,  silver  coin,  red 
litmus  paper,  soda  lime,  porcelain  dish,  kerosene  lamp,  two  bottles. 

(a)  Carbon,  (i)  Recall  or  repeat  the  experiments  which  showed 
that  carbon  is  a  constituent  of  wood,  cotton,  starch,  sugar,  illuminating 
gas  and  candle  wax.  (2)  Heat  a  few  drops  of  turpentine  in  a  porce- 
lain dish,  and  then  set  fire  to  it.  Does  it  contain  carbon?  Hold  a 
bottle  over  the  flame  long  enough  to  collect  any  product,  and  then  test 


Organic  Compounds.  543 

the  contents  for  carbon  dioxide  with  limewater ;  does  the  observation 
verify  the  previous  conclusion?  (3)  Repeat  (2)  with  alcohol.  Does 
alcohol  contain  carbon?  (4)  Burn  a  small  lump  of  camphor  in  a  porce- 
lain dish  or  on  a  block  of  wood.  Does  it  contain  carbon?  (5)  Hold 
a  bottle  over  a  burning  kerosene  lamp  long  enough  to  collect  any  prod- 
uct, and  test  as  in  (2).  Does  kerosene  contain  carbon? 

(b)  Hydrogen,     (i)  Set  fire  to  a  few  drops  of  the  following  liquids 
in  a  porcelain  dish,  and  hold  over  each  flame  a  cold  dry  bottle  long 
enough  to  allow  the  condensation  of  the  water  vapor,  which  is  always 
one  product  of  the  combustion  of  organic  compounds  which  contain 
hydrogen :    alcohol,  turpentine,  kerosene.     (2)  Heat  in  separate  test 
tubes  the  following  dry  solids,  and  if  they  contain  hydrogen,  a  little 
water  vapor  will  condense  on  the  upper  part  of  the  test  tube :  sugar, 
starch,  flour,  wood,  paper,  hair.      (3)  Hold  a  cold,  dry  bottle  for  a  few 
seconds  over  a  burning  kerosene  lamp,  a  Bunsen  flame,  an  ordinary 
gas  flame,  a  burning  candle,  a  burning  taper,  and  describe  the  result. 
Is  hydrogen  a  component  of  kerosene,  illuminating  gas,  and  wax? 

(c)  Oxygen,  which  unites  with  the  hydrogen  of  organic  compounds 
to  form  the  water,  may  come  from  the  compound,  as  in  the  case  of 
sugar,  starch,  wax,  wood,  paper,  or  it  may  come  from  the  air.     No 
simple  experiment  will  determine  the  source  of  the  oxygen. 

(d)  Nitrogen.     Mix  a  little  granulated  gelatine  (one  part)  with  dry 
soda  lime  (two  parts)  and  heat  the  mixture  in  a  test  tube.     Hold  a 
piece  of  moist  red  litmus  paper  in  the  escaping  vapor.     It  will   be 
turned  blue  by  escaping  ammonia  gas.     Gelatine  (also  horn,  glue,  and 
leather)  contains  nitrogen,  which  is  liberated  in  combination  with  hydro- 
gen as  ammonia  gas. 

(^)  Sulphur,  (i)  Put  a  little  mustard  paste  on  a  clean  silver  coin. 
The  brown  stain  is  silver  sulphide.  Explain.  (2)  Why  is  a  silver 
spoon  tarnished  by  a  cooked  egg? 

Draw  a  general  conclusion  regarding  the  composition  of  organic 
compounds. 

Experiment  165.  —  Preparation  and  Properties  of  Alcohol. 
(Teacher's Experiment.)  Materials:  Grape  sugar,  yeast,  limewater, 
bone  black,  sodium  hydroxide.  The  apparatus  consists  of  a  large  bottle 
provided  with  a  one-hole  rubber  stopper  fitted  with  a  delivery  tube 
(like  C,  Z>,  E,  in  Fig.  104)  which  reaches  to  the  bottom  of  a  small 
bottle;  the  latter  has  a  two-hole  stopper.  The  delivery  tube  passes 


544  Experiments. 

through  one  hole,  and  through  the  other  passes  a  bent  tube  connected 
with  a  U-tube. 

I.  Put  a  liter  of  water  in  the  bottle,  add  150  grams  of  grape  sugar, 
and  shake  until  dissolved;  pour  150  cubic  centimeters  of  yeast  into  this 
solution.     Fill  the  small  bottle  half  full  of  limewater.     Fill  the  U-tube 
with  pieces  of  sodium  hydroxide.     Connect  the  apparatus  and  stand  it 
in  a  dark  place,  where  the  temperature  is  25°-3o°  C. 

Fermentation  begins  at  once,  and  carbon  dioxide  —  one  of  the  prod- 
ucts—  bubbles  through  the  limewater,  which  is  protected  from  the 
action  of  the  air  by  the  sodium  hydroxide.  Examine  the  stopper  for  a 
leak,  if  no  change  occurs  in  the  limewater.  The  operation  should  be 
allowed  to  continue  at  least  a  day,  and  longer  if  possible.  The  flask 
will  then  contain  mainly  water,  unchanged  grape  sugar,  alcohol,  and 
some  products  of  minor  importance.  Pour  off  the  liquid,  agitate  it 
with  a  little  bone  black  to  remove  the  odor  and  color,  and  filter.  The 
alcohol,  which  varies  in  quantity  with  the  conditions,  is  dissolved  in  a 
large  excess  of  water  and  must  be  separated  by  distillation. 

II.  The  distillation  is  performed  with  the  apparatus  used  in  Experi- 
ment 13.    Fill  the  flask  half  full  of  the  liquid  from  I,  add  a  few  pieces  of 
pipestem  (or  granulated  zinc,  or  glass  tubing)  to  prevent  "  bumping," 
and  distil  about  50  cubic  centimeters.     Save  the  distillate.     Replace  the 
residue  in  the  flask  by  nrore  liquid  from  I,  distil  again,  and  repeat  this 
operation  until  all  the  liquid  has  been   used.     Replace   the   one-hole 
stopper  with  a  two-hole  stopper,  insert  a  thermometer  in  one  hole  so 
that  the  bulb  just  touches  the  surface  of  the  combined  distillates,  which 
should  now  be  distilled.    Heat  gently,  and  collect  in  a  separate  receiver 
the  distillate  which  is  formed  when  the  liquid  boils  between  80°  and 
93°  C.     This  distillate  contains  most  of  the  alcohol. 

Test  as  follows  :  — 

(a)  Note  the  odor. 

($)  Drop  a  little  into  a  warm  dish,  and  hold  a  lighted  match  over  it. 
If  it  does  not  burn,  it  shows  that  the  alcohol  is  too  dilute.  Put  a  little 
in  a  dish,  warm  gently,  and  light  the  vapor.  Describe  the  result. 

Experiment  166.  —  Properties  of  Alcohol.  (Optional. )  Materials : 
Alcohol,  camphor,  shellac,  rosin,  porcelain  dish. 

(«)  Determine  cautiously  the  odor  and  taste  of  alcohol.  Drop  a  little 
on  a  glass  plate  or  on  a  piece  of  paper,  and  watch  it  evaporate.  Is  its 
rate  of  evaporation  more  rapid  than  that  of  water? 


Organic  Compounds.  545 

(£)  Weigh  a  measured  quantity  (about  25  cubic  centimeters)  of  95 
per  cent  alcohol  and  calculate  its  specific  gravity. 

(c)  Alcohol  dissolves  many  organic  substances.     Try  camphor,  pow- 
dered shellac,  or  rosin.     Describe  the  result.    Verify  the  solvent  power 
of  alcohol  by  adding  water  to  the  solutions.     Describe  the  result. 

(d)  Burn  a  little  alcohol  in  a  dish  and  observe  the  nature  of  the 
flame.     What  are  the  products  of  combustion? 

Experiment  167.  —  Preparation  and  Properties  of   Aldehydes. 

Materials:  Concentrated  hydrochloric  acid,  ethyl  alcohol,  potassium 
dichromate  solution,  methyl  alcohol,  copper  wire,  forceps. 

(#)  Acetic  Aldehyde.  Add  a  little  concentrated  hydrochloric  acid 
and  several  drops  of  ethyl  alcohol  to  several  cubic  centimeters  of  potas- 
sium dichromate  solution.  Warm  gently,  and  observe  the  peculiar- 
smelling  gaseous  product.  It  is  aldehyde  vapor,  aldehyde  itself  being  a 
colorless,  extremely  volatile  liquid,  which  boils  at  20.8°  C. 

(b)  Formic  Aldehyde  or  Formaldehyde.  Put  a  few  cubic  centimeters 
of  methyl  alcohol  in  a  test  tube  and  stand  the  test  tube  in  a  rack.  Wind 
a  piece  of  copper  wire  into  a  spiral  around  a  glass  rod  or  lead  pencil. 
Slip  the  spiral  from  the  rod,  grasp  one  end  into  the  forceps,  and  heat  the 
wire  red-hot  in  the  flame.  Then  quickly  drop  it  in  the  methyl  alcohol. 
The  pungent  vapor  which  is  suddenly  produced  is  largely  the  vapor 
of  formaldehyde. 

Experiment  168.  —  Properties  of  Ether.  Materials :  Ether,  evapo- 
rating dish,  glass  plate,  wax. 

Precaution.  Ether  vapor  is  easily  ignited,  and  should  never  be 
brought  near  a  flame. 

(a)  Pour  a  little  ether  into  a  dish  or  test  tube  and  observe  the  odor 
and  volatility.  Taste  cautiously.  Pour  a  drop  upon  a  glass  plate  or  a 
block  of  wood.  How  does  its  rate  of  evaporation  compare  with  that  of 
alcohol  ?  Pour  a  little  upon  the  hand  and  describe  the  result. 

(£)  Add  a  bit  of  wax  to  a  few  cubic  centimeters  of  ether,  and  shake. 
The  result  is  typical ;  draw  a  conclusion. 

Experiment  169.  —  Properties  of  Acetic  Acid. 

Treat  acetic  acid  as  follows  :  — 

(a)  Taste  (cautiously),  and  describe. 

(b)  Test  with  litmus  paper,  and  describe  the  result. 


546  Experiments. 

(V)  Warm  a  little  in  a  test  tube,  and  smell  (cautiously) .  Describe 
the  odor. 

Experiment  170. — Properties  of  Vinegar. 

(a)    Show,  experimentally,  that  vinegar  contains  acetic  acid. 

(^)    Repeat  Experiment  60,  using  vinegar  instead  of  indigo  solution. 

Experiment  171.  —  Test  for  Acetic  Acid  and  Acetates. 

Cautiously  add  a  few  drops  of  concentrated  sulphuric  acid  to  equal 
(and  small)  volumes  of  acetic  acid  and  alcohol.  Shake  and  warm 
gently.  The  pleasant,  fruitlike  odor  is  due  to  the  vapor  of  ethyl  ace- 
tate, a  volatile  liquid  which  is  always  formed  under  these  circumstances. 

(NOTE.  — This  experiment  is  also  a  test  for  alcohol.) 

Experiment    172.  —  Preparation   and    Properties  of   Acetates. 

Materials  for  (a)  :  Sodium  carbonate,  acetic  acid,  concentrated  sul- 
phuric acid,  alcohol,  porcelain  (or  agate)  dish.  For  (b) :  Litharge, 
acetic  acid,  porcelain  dish. 

Prepare  one  or  both  of  the  following  acetates  :  — 

(a)  Sodium  acetate.     Dissolve  20  grams  of  sodium  carbonate  in  10 
cubic  centimeters  of  water  in  a  porcelain  (or  agate)  dish,  and  slowly 
add  30  cubic  centimeter?  of  commercial  acetic  acid,  with  constant  stir- 
ring.    If  the  solution  is  not  acid,  add  a  little  more  acetic  acid.     Filter 
the  solution,  if  not    clear.     Evaporate   to  crystallization.     When   the 
crystals  have  formed,  remove  and  dry  them.     Describe  the  crystals. 
Prove  that  they  contain  water  of  crystallization.     Test  the  acetate  as 
follows:    (i)    Dissolve  a  little  in  water,  add  a  few  drops  of  concen- 
trated sulphuric  acid,  and  boil.     What  does  the  odor  show  is  present? 
What  other  acids  have  been  similarly  prepared?     (2)    Dissolve  as  in 
(i),  add  a  few  drops  of  alcohol  and  of  sulphuric  acid,  and  boil.     What 
does  the  odor  conclusively  prove?     Preserve  the  crystals,  finally,  in  a 
glass-stoppered  bottle,  or  in  one  having  a  cork  covered  with  paraffin. 

(b)  Lead  acetate  {poisonous).     To  10  grams  of  litharge  add  18  cubic 
centimeters  of  commercial  acetic  acid  in  small  portions.     Stir  the  mix- 
ture constantly  during  the  addition  of  acid.     After  all  the  acid  has  been 
added,  heat  gently  until  the  action  ceases.     (If  the  solution  is  green  or 
bluish,  it  is  due  to  a  copper  compound.     The  copper  may  be  precipi- 
tated and  removed  mechanically  by  standing  a  strip  of  lead  in  the  solu- 
tion for  an  hour  or  more.    Pour  off  the  clearer  liquid  and  then  filter.) 


Organic  Compounds.  547 

Evaporate  cautiously  to  crystallization.  Remove  the  crystals  from 
the  liquid,  and  dry  at  a  moderate  temperature.  Preserve  the  crystals 
finally  as  in  (a).  Describe  the  crystals.  Test  them  for  lead  (see  Exp. 
145  (£)),  and  for  an  acetate. 


Experiment  173.  —  Properties  of  Certain  Organic  Acids.  Mate- 
rials: Tartaric  and  citric  acids,  potassium  permanganate  solution, 
sodium  bicarbonate,  sugar,  concentrated  nitric  acid,  evaporating  dish, 
litmus  paper. 

(1)  Tartaric  acid.     Observe  and  describe  the  results  in  the  follow- 
ing :     (a)  Taste  cautiously  a  dilute  solution  of  tartaric  acid,     (b)  Apply 
the  litmus  test,     (c)  Add  a  little  of  the  solution  to  a  sodium  bicarbonate 
solution.     (W)  Dissolve  two  or  three  crystals  of  potassium  permanga- 
nate in  a  test  tube  half  full  of  water,  add  a  little  sodium  hydroxide  solu- 
tion and  two  or  three  pieces  of  tartaric  acid  (solid).     Warm  gently,  but 
do  not  shake.     The  change  is  due  to  the  reduction  of  the  potassium 
permanganate  by  the  tartaric  acid. 

(2)  Citric  acid.     Proceed  as  in  (i)  with  citric  acid. 

(3)  Oxalic  Acid,     (a)  This  acid  is  poisonous.    Do  not  taste  it. 
(ti)  and  (c)  Proceed  as  in  (i).     (*/)  Dissolve  two  or  three  crystals 
of  potassium  permanganate  in  a  test  tube  half  full  of  water  and  add  half 
the  volume  of  sulphuric  acid.     Add  oxalic  acid  solution  until  a  decided 
change  appears.     Describe  and  explain  it.     (e)  Add  a  few  drops  of 
ink  to  oxalic  acid  solution,  and  shake.     Describe  the  result. 

Experiment  174.  —  Preparation  and  Properties  of  Ethyl  Ace- 
tate. 

Repeat  Experiment  171. 
ANSWER  : 

(1)  What  class  of  organic  compounds  does  ethyl  acetate  represent  ? 
What  general  property  has  this  class  ? 

(2)  To  what  inorganic  compound  does  ethyl  acetate  correspond? 

(3)  What  is  the  relation  of  ethyl  acetate  to  (a)  alcohol  and  (^)  ace- 
tic acid  ? 

Experiment  175.  —  Preparation  of  Soap. 

Prepare  soap  in  an  iron  or  a  tin  dish  by  one  of  the  following 
methods  :  — 


548  Experiments. 

(a)  Dissolve  10  grams  of  sodium  hydroxide  in  75  cubic  centimeters  of 
water,  add  30  grams  of  lard,  and  boil  until  the  mixture  begins  to  solidify. 
Then  add  20  grams  of  fine  salt  in  small  portions.  Stir  constantly  during 
the  addition  of  the  salt.  Boil  a  few  minutes.  Let  the  mass  cool,  and 
then  remove  the  soap,  which  will  form  in  a  cake  at  the  surface. 

(#)  Dissolve  13  to  15  grams  of  sodium  hydroxide  in  100  cubic  centi- 
meters of  water,  add  100  cubic  centimeters  of  castor  oil,  and  boil  for 
about  half  an  hour.  Add  20  grams  of  salt,  and  then  proceed  as  in  (#). 

(c)  Dissolve  8  grams  of  potassium  hydroxide  in  150  cubic  centi- 
meters of  alcohol,  add  10  grams  of  lard,  and  stir  constantly  while  the 
mixture  is  being  heated  cautiously  to  sirupy  consistency.  Allow  the 
solution  to  cool.  The  jellylike  product  is  soap. 

Preserve  a  sample. 

Experiment  176. —  Properties  of  Soap.  Materials:  Soap,  sul- 
phuric acid,  calcium  sulphate,  magnesium  sulphate,  and  acid  calcium 
carbonate  solutions. 

Test  as  follows  the  soap  prepared  in  Experiment  175  :  — 

(#)  Leave  soap  shavings  exposed  to  the  air  for  several  days.  What 
does  the  result  show  about  the  presence  of  water  in  the  soap  ? 

(£)    Test  soap  solution  with  litmus  paper. 

(c)  Add  considerable  dilute  sulphuric  acid  to  a  soap  solution.     The 
precipitate  is  a  mixture  mainly  of  palmitic  and  stearic  acids.    Describe  it. 

(d)  To  a  little  soap  solution  in  separate  test  tubes  add  calcium  sul- 
phate and  magnesium  sulphate  solutions.     Describe  the  result.     Boil 
for  a  few  minutes  and  describe  the  result.     Prepare  a  solution  of  acid 
calcium  carbonate  by  passing  carbon  dioxide  into  limewater  until  the 
precipitate  is  redissolved  (see  Exp.  69).     Add  some  of  the  solution  to 
a  soap  solution,  and  describe  the  result.     Boil,  as  above,  and  describe 
the  result. 

ANSWER  : 

(1)  -What  is  hard  water  ?     Soft  water  ? 

(2)  What   is   permanent   hardness?     Temporary   hardness?     How 
can  the  later  be  removed  ? 

Experiment  177.  —Properties  of  Glycerine. 

O)  Add  a  little  glycerine  to  a  test  tube  half  full  of  water,  and  shake. 
Add  considerable  more  glycerine,  and  shake.  What  does  the  result 
show  about  the  solubility  of  glycerine  in  water? 


Laboratory  Equipment.  549 

(£)  Cautiously  taste  the  liquid  resulting  from  (a).  Describe  the 
result. 

Experiment  178.  —  Fehling's  Test  for  Sugar.  Materials:  Copper 
sulphate,  Rochelle  salt,  sodium  hydroxide,  and  grape  sugar  solutions. 

Mix  equal  (and  small)  volumes  of  copper  sulphate,  Rochelle  salt,  and 
sodium  hydroxide  solutions  in  a  test  tube,  and  boil  carefully.  The 
mixture  should  be  strongly  alkaline.  Add  a  little  grape  sugar  solution, 
and  boil  until  a  decided  change  is  produced.  The  precipitate  is  cuprous 
oxide.  Describe  it. 

(NOTE.  —  Cane  sugar  must  be  changed  to  grape  sugar  by  boiling 
with  dilute  sulphuric  acid  before  the  above  test  is  applicable.) 

Exercises  for  Review. 

1.  What  happens  to  sugar  and  starch   (a)  when  heated,  and   (£) 
when  treated  with  concentrated  sulphuric  acid  ? 

2.  What  is  the  test  for  starch  ? 

3.  Discuss  the  solubility  of  alcohol  in  water. 

4.  What  is  the  effect  of  heat  upon  paper  and  cotton?     Of  potassium 
permanganate  on  paper? 

Experiment  179.  — Properties  of  Benzene. 

Put  one  or  two  drops  of  benzene  in  an  evaporating  dish,  and 
cautiously  bring  a  lighted  match  near  it.  Describe  the  result. 


LABORATORY     EQUIPMENT. 

The  Equipment  of  a  laboratory  should  be  limited  solely  by  the  means 
at  the  disposal  of  the  teacher.  Accurate  and  rapid  work  is  largely 
determined  by  the  available  facilities,  and  no  pains  should  be  spared 
to  secure  the  equipment  which  will  yield  the  largest  educational' return 
for  the  time  and  money  expended. 

The  lists  given  below  include  the  apparatus  and  chemicals  needed  for 
the  experiments  in  this  book.  Quantities  and  prices  have  been  omitted 
in  justice  to  teachers,  dealers,  and  the  author.  Different  teachers  use 
different  quantities,  prices  fluctuate,  and  qualities  vary.  The  author,  at 
his  own  suggestion,  has  lodged  with  the  L.  E.  Knott  Apparatus  Co.,  16 
Ashburton  Place,  Boston,  Mass.,  information  regarding  the  quantities 


550 


Experiments. 


of  apparatus  and  chemicals  used  by  his  classes.  It  is  hoped  that  teach- 
ers will  correspond  with  both  author  and  dealer  when  preparing  order 
lists.  The  author  takes  this  opportunity  to  say  that  he  has  no  financial 
connection  whatever  with  any  dealer  in  scientific  supplies. 


LIST  A.  — INDIVIDUAL  APPARATUS. 

This  list  includes  the  apparatus  constantly  used  by  a  single  student, 
who  should  be  provided  with  each  piece.  The  set  will  cost  from  $4.75 
to  $5.  The  discount  on  apparatus  in  this  and  succeeding  lists  depends 
upon  the  total  amount  of  the  order. 


6  Test  tubes,  6  x  |. 
3  Test  tubes,  8  x  i . 

Test-tube  holder. 

Test-tube  rack. 

Test-tube  brush. 

Bunsen  burner. 

Blowpipe. 

Blowpipe  tube. 

Bottles,  wide  mouth,  250  cc. 

Funnel,  2.\  in. 

Evaporating  dish. 

Pair  iron  forceps. 

Triangular  file. 


i  Mortar  and  pestle,  3  in. 
i  Deflagrating  spoon. 

1  Pneumatic  trough. 

2  ft.  Rubber  tubing,  \  in.  in  diam. 
100  Filter  papers,  4  in. 

i  ft.  Glass  rod. 

6  in.  Rubber  tubing,  TV 

i  One-hole  and  i  two-hole  rubber 

stopper  to  fit  large  test  tube. 
4  ft.    Glass    tubing    to   fit  rubber 

stoppers  (above), 
i  Safety  tube. 


LIST  B.  —  SPECIAL  APPARATUS. 

This  list  includes  apparatus  used  occasionally.     Numbers  in  paren- 
theses refer  to  experiments.     The  set  will  cost  from  $3  to  $3.25. 


i  Crucible,  Hessian,  4  in.  deep  (59, 
90). 

Dish,  lead  (80). 

Flask,  Erlenmeyer,  250  cc.  (25). 
Pinchcock,  Mohr  (96). 
Screw,  Hofmann  (25). 
Thistle  tube  (96) . 
Lamp  chimney  (77). 


i  Graduated   cylinder,   25   cc 

and  others), 
i  Magnet  (156). 
i  Candle  (63,  77). 
i  Sand-bath  pan,  4  in. 
i  Wing-top  burner  (Int.  § 
i  Dish,  iron  or  tin  (109,  in) 
i  Retort,  250  cc.  (49). 


(25 


List  E.  —  Chemicals. 


551 


LIST  C.  — APPARATUS    FOR  TEACHER'S   EXPERIMENTS. 

This  list  includes  the  additional  apparatus  for  the  Teacher's  Experi- 
ments. Numbers  in  parentheses  refer  to  experiments.  The  set  will 
cost  about  $11. 


Electrolysis  apparatus  (22,  36). 
Flask,  500  c.c.  (10,  13,  165). 
Two-hole  rubber  stopper  for  above. 
U-tube  (10,  73,  165). 

2  One-hole    rubber    stoppers    for 
above. 

4  in.  Capillary  tubing  (10,  22,  36). 

3  ft.  Glass    tubing    to    fit   rubber 
stoppers. 


i  Safety  tube  (10). 

i  Condenser  complete  (13,  165). 

i  Tripod  (13). 

i  Thermometer  (165). 

i  Chlorine  tube  (23). 

i  Ignition  tube,  6  in.  (73). 

i  Bottle,  wide  mouth,  50  cc.  (10). 

Battery,  3  cells  (Grenet)  (22, 36, 37). 

i  Bottle,  2000  cc.  (165). 


LIST  D.—  GENERAL   APPARATUS. 

This  list  includes  the  general  laboratory  apparatus.  It  should  be 
extended  as  demands  arise.  It  does  not  include  such  items  as  dupli- 
cate stoppers,  extra  glassware,  tools,  etc.  Special  inexpensive  articles 
are  noted  in  the  experiments  and  in  the  "Handbook  for  Teachers" 
accompanying  this  book. 


Corks,  assorted. 

Copper  wire,  No.  24. 

Glass  plates,  4  x  4  in. 

Iron  stands,  3  rings,  2  clamps. 

Matches. 

Wire  gauze,  iron,  4  x  4  in. 

Wooden     blocks,    6  x  6  x  i 


in., 


6  x  6  x  f  in.,  4  x  4  x  £  in.  (with 
I-  in.  hole  in  center  —  see  Exp. 
38). 


Sand. 

Wood,    thin    sticks    (Exp.   6   and 

others) . 

Rule,  foot  and  30  cm. 
Scales,  trip. 
Weights  for  above. 
Tapers. 
Emery  paper. 
Kerosene  lamp. 
Graduated  cylinders,  500  cc.,  100  cc. 


LIST   E.  —  CHEMICALS. 

This  list  includes  the  chemicals  needed  for  this  book.     Numbers  in 
parentheses  refer  to  experiments  in  which  the  chemicals  are  used. 


Acid,  acetic, 
citric. 


Acid,  hydrochloric, 
nitric. 


552 


Experiments. 


Acid,  oxalic. 

pyrogallic  (25). 
sulphuric, 
tartaric. 
Alcohol,  ethyl. 

methyl  (167). 
Alum,  chrome. 

potassium. 
Aluminium,  metal. 

sulphate. 

Ammonium,  chloride, 
hydroxide, 
nitrate, 
oxalate. 
sulphide. 

Arsenious  oxide  (95,  105). 
Asbestos,  shredded  (73). 
Baking  powder  (68) . 
Barium  chloride. 

nitrate  (125). 
Benzene  (179). 
Bismuth  (107). 
Bleaching  powder. 
Borax  (powd.). 
Cadmium  chloride  (131). 
Calcium  carbide  (72). 

carbonate  (marble), 
chloride, 
fluoride  (80). 
oxide  (lime), 
sulphate. 

Carbon  disulphide. 
Chalk  (native)  (68). 
Charcoal,  animal  (powd.). 
lump. 

wood  (powd.). 
Coal,  soft. 
Cobalt  nitrate. 
Cochineal. 


Coin  (silver). 
Copper  nitrate, 
sheet. 

sulphate  (cryst.). 
Cotton  (absorbent). 
Cream  of  tartar. 
Ether. 

Galena  (148). 
Gelatine. 
Glycerine. 
Gold  leaf  (book). 
Hematite  (161). 
Indigo. 
Iodine. 
Iron,  chloride  (#:). 

filings. 

powder. 

pyrites  (161). 

sulphate  (ous). 

sulphide  (ous). 

wire  (fine). 

wrought. 
Kerosene. 
Lead  acetate. 

carbonate. 

dioxide  (peroxide). 

nitrate. 

monoxide  (litharge). 

sheet. 

tea. 

tetroxide. 
Limonite  (161). 
Litmus  paper. 
Magnesium  oxide, 
ribbon, 
sulphate. 
Magnetite  (161). 
Manganese  dioxide, 
sulphate. 


Solutions. 


553 


Mercury. 

Mercuric  chloride, 
nitrate, 
oxide  (7). 
Mercurous  nitrate. 
Mustard. 

Nickel  chloride  (162). 
Paraffin. 

Phenolphthalein  (66). 
Picture  cord  (iron)  (6). 
Potassium,  metal  (no). 

bromide. 

carbonate. 

chlorate  (cryst.). 

chlorate  (powd.). 

chloride. 

chromate. 

dichromate. 

ferricyanide. 

ferrocyanide. 

hydroxide. 

iodide. 

nitrate. 

permanganate. 

sulphate. 

sulphocyanide  (thiocya- 

nate)  (157-158). 
Pyrite  (161). 
Rochelle  salt  (178). 
Rosin. 

Selenite  (gypsum,  cryst.). 
Shellac. 
Siderite  (161). 
Silver  nitrate. 
Soap. 


Soda  lime  (164). 
Sodium,  metal. 

bicarbonate. 

carbonate. 

chloride. 

hydroxide. 

hyposulphite  (thiosul- 
phate). 

nitrate. 

phosphate   (disodium 
phosphate) . 

silicate  (101). 

sulphate. 

sulphite  (96). 
Stannous  chloride  (tin  crystals). 
Starch. 
Steel. 

Strontium  nitrate  (123). 
Sugar,  cane. 

grape  (165). 
Sulphur,  flowers. 

roll. 

Tartar  emetic. 
Tin,  granulated. 
Tooth  powder  (68). 
Turpentine. 
Vaseline. 
Vinegar. 
Water,  distilled. 
Whiting  (68). 
Wood  ashes. 
Zinc,  granulated, 
oxide, 
sheet, 
sulphate. 


SOLUTIONS. 

The  following  solutions  are  needed  .for  the  experiments  in  this  book. 
Those  not  included  are  described  in  the  experiments  requiring  their  use. 


554 


Experiments. 


Alum,  10  per  cent. 

Ammonium  chloride,  10  per  cent. 

Ammonium  hydroxide,  i  vol.  to  3 
vols.  water. 

Ammonium  oxalate,1  4  per  cent. 

Ammonium  sulphide,  i  vol.  to 
i  vol.  water. 

Barium  chloride,2  5  per  cent. 

Battery  solution  (Grenet).  Dis- 
solve 103  gm.  powdered  potas- 
sium dichromate  in  i  liter  of 
water  and  slowly  add  103  gm. 
cone,  sulphuric  acid  with  con- 
stant stirring. 

Calcium  chloride,  10  per  cent. 

Chlorine  water,1  saturated  (see 
Exp.  23,  38). 

Cobalt  nitrate,  5  per  cent. 

Cochineal.  Prepare  as  described 
under  Indigo. 

Copper  sulphate,  10  per  cent. 

Disodium  phosphate,  lo.per  cent. 

Ferric  chloride,  5  per  cent. 

Ferrous  sulphate,1  10  per  cent. 

Hydrochloric  acid,  i  vol.  to  4  vols. 

Indigo.  Grind  a  little  with  water 
and  dilute  as  desired. 

Iodine.  Grind  to  solution  12  gm. 
iodine,  20  gm.  potassium  iodide, 
10  cc.  water,  and  add  to  1000  cc. 
water. 

Lead  acetate,  10  per  cent. 
Lead  nitrate,  10  per  cent. 
Limewater.      Let  water  stand  over 
lime  for  several  days,  and  siphon 
off  the  clear  liquid. 


Magnesium  sulphate,  10  per  cent. 

Manganese  chloride,  10  per  cent. 

Mercuric  chloride,  5  per  cent.  Poi- 
son. 

Mercurous  nitrate,8  5  per  cent. 

Nitric  acid,  i  vol.  to  4  vols.  water. 

Potassium  bromide,  5  per  cent. 

Potassium  chloride,  5  per  cent. 

Potassium  chromate,  10  per  cent. 

Potassium  dichromate  (or  bichro- 
mate), 5  per  cent. 

Potassium  ferricyanide,  10  per 
cent. 

Potassium  ferrocyanide,  10  per 
cent. 

Potassium  hydroxide,  10  per  cent. 

Potassium  iodide,  5  per  cent. 

Potassium  nitrate,  10  per  cent. 

Potassium  permanganate,2  5  per 
cent. 

Potassium  sulphate,  10  per  cent. 

Potassium  thiocyanate  (or  sulpho- 
cyanide),  i  per  cent. 

Silver  nitrate,  5  per  cent. 

Sodium  carbonate,  10  per  cent. 

Sodium  chloride,  10  per  cent. 

Sodium  hydroxide,  10  per  cent. 

Stannous  chloride.1  Dissolve  500 
gm.  of  the  salt  in  1000  cc.  hot 
cone,  hydrochloric  acid,  and  add 
a  piece  of  tin. 

Sulphuric  acid,  i  vol.  to  4  vols. 
water. 

Tartar  emetic,  10  per  cent. 

Zinc  sulphate,  10  per  cent. 


1  Must  be  freshly  prepared.  2  Use  distilled  water. 

8  Use  distilled  water,  and  add  75  cc.  concentrated  nitric  acid  and  a  little  mercury. 


INDEX. 


Absolute  zero,  439. 
Acetates,  417. 

Ethyl,  419. 

Metallic,  419. 
Acetic  acid,  415. 

Constitution,  170,  416. 

Glacial,  415. 

Preparation,  415. 

Properties,  415. 

Series,  414. 

Test,  419. 
Acetylene,  116,  205. 

As  illuminant,  206. 

Burner,  207. 

Composition,  206. 

Explosive  properties,  205. 

Flame,  207,  216. 

Generation,  207. 

Liquid,  205. 

Series,  202. 
Acetylides,  206. 
Acid,  acetic,  415. 

Benzoic,  431. 

Boracic,  261. 

Boric,  261. 

Butyric,  417. 

Capric,  421. 

Caproic,  421. 

Carbolic,  431. 

Carbonic,  194. 

Chloric,  91. 

Chlorous,  91. 

Citric,  419. 

Cyanic,  198. 

Ethyl  sulphuric,  414. 

Fuming  sulphuric,  251. 


Acid,  continued. 
Gallic,  432. 
Glacial  acetic,  415. 
Glacial  phosphoric,  268. 
Hydriodic,  232. 
Hydrobromic,  230. 
Hydrochloric,  140. 
Hydrocyanic,  198. 
Hydrofluoric,  227. 
Lactic,  418.    — • — 
Malic,  418. 
Metaphosphoric,  268. 
Metastannic,  355. 
Muriatic,  92,  140. 
Nitric,  154. 
Nitrosylsulphuric,  248. 
Nitrous,  159. 

Nordhausen  sulphuric,  252. 
Orthophosphoric,  268. 
Oxalic,  417. 
Palmitic,  417. 
Perchloric,  91. 
Picric,  431. 
Prussic,  198. 
Pyrogallic,  431. 
Pyroligneous,  415. 
Pyrophosphoric,  269. 
Pyrosulphuric,  252. 
Salicylic,  431. 
Silicic,  257. 
Stearic,  417. 
Sulphocyanic,  198. 
Sulphuric,  246. 
Sulphurous,  244,  245. 
Tannic,  432. 
Tartaric,  418. 


555 


556 


Index. 


Acid  calcium  sulphate,  245. 

Of  air,  196. 

Oxide,  97. 

Phosphate,  90,  269. 

Potassium  fluoride,  226. 

Reaction,  90. 

Salt,  96. 

Sodium  carbonate,  289. 

Sodium  sulphate,  245. 

Sulphates,  251. 
Acidity,  94. 
Acids,  90. 

And  ionization,  129. 

And  oxygen,  18. 

Chlorine,  91. 

Commercial  names,  92. 

Defined,  90. 

Dibasic,  92. 

General  properties,  87. 

In  familiar  substances,  90. 

Monobasic,  92. 

Nomenclature,  91. 

Organic,  92,  414. 

Oxygen  in,  97. 

Relation  of  oxides  to,  96. 

Tribasic,  92. 
Addition  products,  204. 
Agate,  255. 
Air,  61.  4f 

Acid  of,  196. 

Alkaline,  149. 

Bad,  67. 

Composition,  64. 

Dephlogisticated,  18. 

Empyreal,  18. 

Fixed,  196. 

Gravimetric  composition,  66. 

Hydrogen  dioxide  in,  60. 

Liquid,  69. 

Marine  acid,  140. 

Mixture,  69. 

Relative  humidity,  66. 

See  Atmosphere. 

Slaked  lime,  324. 

Solubility,  69. 


Air,  continued. 

Volumetric  composition,  64. 

Weight  of  liter,  62. 
Alabaster,  326. 
Alchemists,  154,  157,   158,  235,  246 

251,  308,  314,  354,  3  7°- 
Alcohol,  ethyl,  409. 

Absolute,  410. 

Commercial,  410,  411. 

Constitution,  407. 

Fermentation,  416. 

Formulas,  407. 

Oxidation,  412. 

Preparation,  410. 

Pure,  410. 

Test,  419. 

Uses,  410. 
Alcohol,  methyl,  409. 

Triacid,  420. 

Wood,  409. 
Alcoholic  liquors,  411. 
Alcohols,  408. 

Constitution,  409. 

General  nature,  408. 
Aldehyde,  acetic,  412. 

Benzoic,  431. 

Formic,  412. 

Salicylic,  432. 
Aldehydes,  412. 
Alizirin,  367,  432. 
Alkali,  92,  93,  94. 

Action  on  litmus,  92. 

And  glass,  259. 

Metals,  284. 

Sodium  carbonate,  289. 

Volatile,  93,  149. 
Alkalies,  common  names,  92. 

Fixed,  93. 

Properties,  93. 
Alkaline,  92. 

Air,  149. 

Earth  metals,  319. 

Reaction,  92. 

Silicate,  257,  258. 
Alkaloids,  433. 


Index. 


557 


Allotrope,  191. 

Allotropic  modification,  191. 

Carbon,  190. 

Silicon,  255. 

Sulphur,  239. 
Allotropism,  190. 
Allotropy,  191. 
Alloys,  282. 

Antimony,  356,  360. 

Copper,  305. 

Fusible,  337,  360. 

Lead,  360. 

Magnesium,  346. 

Manganese,  369. 

Mercury,  339. 

Nickel,  306. 

Platinum,  394. 

Silver,  311. 

Tin,  356. 

Zinc,  306,  336. 
Allylene,  202. 

dels,  338. 

rn,  349. 

nium,  349. 
349. 

Cake,  348. 

Chrome,  350,  367,  368. 

Concentrated,  349. 

General  formula,  350. 

History,  350. 

Iron,  386. 

Potassium,  349. 

Shale,  350. 

Slate,  350. 

Sodium,  349. 
Alumen,  343. 
Alumina,  346. 

Preparation,  347. 

See  Aluminium  oxide. 
Aluminates,  348. 
Aluminium,  343. 

Acetate,  350,  417. 

Alloys,  344,  346. 

Bronze,  305,  346. 

Carbide,  203. 


Aluminium,  continued. 

Chloride,  351. 

History,  343,  344,  345. 

Hydroxide,  348. 

Impurities,  345. 

In  gems,-  347. 

Leaf,  346. 

Metallurgy,  343,  344. 

Name,  343. 

Occurrence,  343. 

Older  processes,  344. 

Oxide,  343,  346,  347. 

Price,  345. 

Production,  344. 

Properties,  345. 

Silicate,  351. 

Sulphate,  348. 

Test,  347. 

Uses,  346. 

Alumino  ferric  cake,  348. 
Aluminum.     See  Aluminium. 
Alumium,  343. 
Alunite,  350. 
Amalgamated  zinc,  339. 
Amalgamation,  281. 

Process  for  silver,  309. 
Amalgams,  defined,  282,  339. 

4fctd»  339- 

Tin,  339,  356. 
Amethyst,  255. 

Oriental,  347. 
Ammonia,  147. 

Anhydrous,  148. 

As  a  refrigerant,  152. 

Composition,  153. 

Formation,  147. 

From  coal,  147,  148. 

In  ice-making,  152. 

Liquefied,  148,  149,  153. 

Muriate  of,  151. 

Near  stables,  147. 

Of  commerce,  148. 

Preparation,  147. 

Properties,  148. 

Soda  process,  289. 


558 


Index. 


Ammonia,  continued. 

Uses,  152. 

Water,  149. 

Ammoniacal  liquor,  213. 
Ammonium,  150. 

Alum,  349. 

Carbonate,  152. 

Chloride,  151. 

Chloroplatinate,  394. 

Compounds,  152. 

Hydroxide,  148,  149,  150. 

Molybdate,  369. 

Nitrate,  152. 

Sulphate,  151. 

Sulphide,  152. 

Sulphocyanate,  152. 

Salts,  150. 
Amorphous,  184. 

Carbon,  181,  184,  190. 

Sulphur,  239,  240. 
Amygdalin,  432. 
Amyl  acetate,  419. 

Valerate,  419. 
Anaesthetic,  412,  413.        » 
Analysis,  3,  50. 

Qualitative,  242. 

Spectrum,  403. 

Water,  39. 
Anglesite,  357. 
Anhydride,  97. 

Carbonic,  195. 

Nitric,  163. 

Nitrous,  163. 
Anhydrite,  326. 
Anhydrous,  46. 
Aniline,  431. 
Dyes,  431. 

Animal  charcoal,  189. 
Anion,  120,  121. 
Annealing  glass,  260. 
Anode,   120,  121,  285,  291,  303,  312, 

332,  344- 
Anthracene,  432. 
Anthracite  coal,  185,  186. 
Antichlor,  138,  245,  252. 


Antidote  for  arsenic  poisoning,  273, 

385. 

Antifriction  metals,  336. 
Antimony,  274. 
Acids,  -274. 
Alloys,  356,  360. 

As  metalloid,  278. 

Chlorides,  275. 

Compounds,  419. 

Name,  274. 

Oxides,  274. 

Oxychlorides,  275. 

Test,  275. 

Trisulphide,  274. 
Apatite,  225,  265. 
Aqua  ammonia,  148,  149. 

Fortis,  158. 

Regia,  160,  316. 
Argentiferous  lead,  308. 
Argentite,  308. 
Argentum,  308. 
Argol,  418. 
Argon,  68,  404. 
Aristotle,  61. 
Armor  plate,  380,  389. 
Arrhenius,  128,  442. 
Arsenic,  272. 

Acids  and  salts,  273. 

Antidote,  273,  385. 

As  metalloid,  278. 

Marsh's  test,  273. 

Ores,  272. 

Oxide,  272. 

Poisoning,  273,  385. 

Production,  272. 

Pyrites,  272. 

Sulphide,  273. 

Test,  273. 

Trioxide,  272. 

Uses,  272. 

Vapor  density,  169. 

White,  272. 
Arsenious  oxide,  272. 
Arsenolite,  272. 
Arsine,  273. 


Index. 


559 


Artificial  diamonds,  346. 

Graphite,  118. 

Stone,  333. 
Asbestos,  331. 
Ash,  black,  288. 

Seaweed,  230,  231. 
Ashes     and     potassium     compounds, 

298. 
Atmosphere,  61. 

An,  62. 

And  plants,  194. 

Argon  in,  68. 

Carbon  dioxide  in,  67. 

Composition,  64. 

Inert  gases  in,  69. 

Ingredients,  62. 

Nitrogen  in,  72. 

Of  sun,  23. 

Oxygen  and  nitrogen  in,  63. 

Ozone  in,  22. 

Pressure,  62. 

Properties,  61. 

Water  vapor  in,  31,  66. 

See  Air. 

Atomic  theory,  79. 
Atomic  weights,  101. 

And  symbols,  103. 

And  valence,  178. 

Classification  by,  397. 

Determination,  170. 

Methods  of  determining,  173. 

Relation  of  properties  to,  398. 

Standards,  102. 

Table,  448. 
Atoms,  79,  81. 

And  ions,  125. 

And  molecules,  80. 

Combining  power,  176,  177,  406. 

In  a  molecule,  168,  174,  191,  204, 
232,  238,  267,  272,  286,  336, 
339,  407,  430. 

Replacement  of,  176. 

Space  relations,  178. 
Attraction,  chemical,  4. 
Aurum,  314. 


Avogadro,  167,  442. 
Hypothesis,  167. 
Azote,  63. 
Azurite,  301,  308. 

Babbitt's  metal,  336. 

Bacteria,  155. 

Baking  powder,  290,  418. 

Soda,  290. 
Balard,  230,  442. 
Bamboo,  257. 
Barite,  329. 
Barium,  328. 

Carbonate,  329. 

Chloride,  329. 

Compounds,  329. 

Dioxide,  59. 

Nitrate,  329. 

Oxides,  12,  59,  329. 

Sulphate,  329,  362. 

Sulphide,  329. 

Test,  329. 
Barley  sugar,  424. 
Baryta  water,  329. 
Base,  92,  93. 

Ammonium  hydroxide  as,  150. 

Diacid,  94. 

Monacid,  94. 

Triacid,  94. 
Bases,  88. 

And  ionization,  129. 

Nomenclature,  93. 

Relation  of  oxides  to,  96. 
Basic,  93. 

Bismuth  nitrate,  276. 

Oxides,  97. 

Salt,  96. 
Basicity,  92. 
Basil  Valentine,  246. 
Bath  metal,  305. 
Battery,  electric,  120. 

Leclanche,  151. 
Baux,  deposits  at,  348. 
Bauxite,  348. 
Becher,  16,  442. 


560 


Index. 


Beef  fat,  420. 
Beehive  oven,  189. 
Beer,  193,  411. 
Beet  sugar,  424. 

Potassium  carbonate  from,  297. 
Bell  metal,  306.        „ 
Bench  of  retorts,  210. 
Benzene,  213,  430. 

Constitution,  430. 

Derivatives,  430. 

Series,  202. 

Source,  429. 
Benzine,  208,  430. 
Benzoic  acid,  431. 

Aldehyde,  431. 
Benzol,  430. 
Bergman,  196,  443. 
Berlin  blue,  388. 
Berthollet,  77,  443. 
Beryllium,  401. 
Berzelius,  443. 

And  Dulong,  57. 
Bessemer,  443. 
Bessemer  steel,  381. 
Beverages,  sparkling,  193. . 
Bicarbonate,  sodium,  195. 
Binary  compounds,  95,  176. 
Bismite,  275. 
Bismuth,  275. 

Carbonate,  275. 

Dioxide,  276. 

Hydroxide,  276. 

Nitrate,  276. 

Oxychloride,  276. 

Pentoxide,  275. 

Subnitrate,  276. 

Sulphide,  275. 

Test,  276. 

Trichloride*  276. 

Trioxide,  275. 
Bismuthinite,  275. 
Bismutite,  275. 
Bisulphite  of  soda,  245. 
Bittern,  229. 
Bituminous  coal,  185,  186,  210, 


Bivalent  elements,  176. 
Black,  196,  334,  443. 
Black  ash  process,  288. 

Damp,  203. 

Lead,  183,  357,  359. 

Magnesia,  370. 

Oxide  of  manganese,  370. 
Blast  furnace,  275,  281. 

Lamp,  29. 
Bleaching  by  chlorine,  136,  137,  138. 

Hydrogen  dioxide,  60. 

Sodium  peroxide,  293. 

Sulphur  dioxide,  244. 
Bleaching  powder,  137. 
Block  tin,  355. 

Pipe,  40. 
Blood  and  oxygen,  16,  17. 

Iron,  373. 

Blow,  water  gas,  213. 
Blowpipe,  222. 

Flame,  29,  222. 

Oxyhydrogen,  17,  28. 
Blue  paint,  417. 

Print  paper,  388. 

Stone,  307. 

Vitriol,  307. 
Bonds,  407. 
Bone  ash,  phosphorus  from,  265. 

Cupel,  310. 
Bone  black,  189. 
Bones,  271. 

Phosphorus  from,  265. 
Books,  reference,  450. 
Boracic  acid,  261. 
Boracite,  261. 
Borax,  261,  262. 

And  soldering,  263. 

Bead,  262. 
Boric  acid,  261. 
Borides,  261. 
Bornite,  301,  373. 
Boron,  261. 
Bort,  182. 
Boyle,  64,  443. 

Law,  19. 


Index. 


S6i 


Brand,  265. 
Brandy,  411. 
Brass,  305. 

Cyprian,  301. 
Braunite,  369. 
Bread  making,  419,  427. 
Breathing,  16,  17. 
Bricks,  352. 
Brimstone,  238. 
Brines,  287. 

Britannia  metal,  306,  356. 
British  coal  fields,  187. 

Gum,  427. 
Brittle  metals,  279. 
Bromides,  230. 
Bromine,  228. 

Commercial  process,  229. 

Compounds,  230. 

Discovery,  230. 

Name,  229. 

Production,  230. 

Properties,  229. 

Uses,  230. 

Water,  229. 
Bronze,  305. 

Aluminium,  305. 

Phosphor,  305. 

Silicon,  305. 
Brown  iron  ore,  374. 
Bullets,  360. 
Bunsen,  64,  219,  284,  403,  443. 

Burner,  219. 

Flame,  219,  220,  221. 
Burette,  98. 
Burner,  acetylene,  207. 

Bunsen,  219. 

Self-lighting,  26. 
Burning,  15. 
Burnt  alum,  349. 
Butane,  203. 
Butter,  421. 

Artificial,  421. 

Rancid,  417. 
Butylene,  204. 
Butyric  acid,  417. 


Cadmium,  337. 

Sulphide,  337. 

Test,  337. 

Vapor  density,  169. 
Caesium,  284,  299,  403. 
Caffeine,  433. 
Cake,  alum,  348. 

Alumino  ferric,  348. 

Press,  296. 

Salt,  288. 
Calamine,  334. 
Calcarone,  236. 
Calcination  of  ores,  281. 
Calcite,  320. 
Calcium,  319. 

Preparation,  319. 

Properties,  319,  320. 

Test,  328. 
Calcium  and  carbonate,  195. 

Acid  sulphate,  429. 

Borate,  262. 

Carbide,  116,  205,  207. 

Carbonate,  192,  195,  319,  320,  321. 

Chloride,  67,  327,  328. 

Fluoride,  225,  227. 

Hydroxide,  325.     See  Limewater. 

Hypochlorite,  137. 

Iodide,  319. 

Light,  29. 

Magnesium  carbonate,  331. 

Manganite,  370. 

Nitrate,  155. 

Oxide,  324.     See  Lime. 

Sulphate,  326,  327. 

Sulphide,  288,  328,  329. 
Calculations,  chemical,  103. 
Calico  printing,  350. 
Calomel,  340. 
Caloric,  112. 
Candle  flame,  216,  217,  218. 

Power,  216. 
Candles,  stearin,  422. 

Illuminating  gas,  216. 
Cane  sugar,  423. 

See  Sugar. 


562 


Index. 


Cannizzaro,  443. 
Capric  acid,  421. 
Caproic  acid,  421. 
Caramel,  424. 
Carat,  diamond,  183. 

Gold,  316. 
Carbide,  aluminium,  203. 

Calcium,  116,  205,  207. 

Iron,  209. 
Carbohydrate,  423. 
Carbolic  acid,  431. 
Carbon,  181. 

Amorphous,  181,  184,  190. 

Boride,  261. 

Combining  power,  202. 

Compounds,  181,  405,  406. 

Bisulphide,  112,  252. 

Gas,  190,  213. 

Silicide,  117,  258. 

Test,  189. 
Carbonado,  182. 
Carbonate,  acid,  195. 

Ammonium,  152. 
Carbonates,  195. 

Normal,  195. 
Carbon  dioxide,  191. 

And  combustion,  193. 

Composition,  195. 

.Detection,  67,  68. 

Formation,  191. 

History,  196. 

In  air,  194. 

In  atmosphere,  67. 

Liquid,  193. 

Occurrence,  191. 

Other  names,  203. 

Preparation,  192. 

Properties,  193. 

Relation  to  life,  194. 

Solid,  193. 

Solubility,  42,  193,  194. 

Test,  192,  325. 
Carbonic  acid,  97,  194. 

Anhydride,  195. 

Oxide,  197. 


Carbon  monoxide,  196,  197. 

In  water  gas,  215. 
Carbonyl  chloride,  198. 

Nickel,  198. 
Carborundum,  117. 

Furnace,  117,  118. 
Carboxyl,  414. 
Carbureter,  213. 
Carbureting,  215. 
Carlisle  and  Nicholson,  53, 
Carnallite,  294,  295,  331. 

Magnesium  from,  332. 
Carnelian,  255. 
Cassiterite,  354. 
Casting  iron,  378. 
Cast  iron,  378. 

Varieties,  378. 
Castner,  284,  285. 
Catalysis,  250. 
Catalytic  action,  250. 
Catalyzer,  250. 
Cathode,  120,  121. 
Cation,  120,  121. 
Caustic  lime,  324. 

Lunar,  312. 

Potash,  297. 

Soda,  290. 
Cavendish,  16,  27,  30,  55,  64,  69,  157, 

443- 

Celestite,  328. 
Cell,  electrolytic,  122. 

Galvanic,  119. 

Voltaic,  119. 
Celluloid,  429. 
Cellulose,  428. 

Nitrates,  428. 
Cementation  process,  380. 
Cements,  325. 

Centigrade  thermometer,  439. 
Cerussite,  357,  361. 
Chalcedony,  255. 
Chalcopyrite,  301,  373. 
Chalcorite,  301. 
Chalk,  322,  323. 
Chalybeate  water,  37,  387. 


Index. 


563 


Champagne,  193. 
Changes,  I. 
Chaptal,  72. 
Charcoal,  187. 

Animal,  187. 

Pit,  1 88. 

Wood,  187. 
Charles,  444. 

Law,  19. 

Checkerberry,  432. 
Chemical  action,  3,  ill,  250. 

And  electricity,  119. 

And  heat,  112. 

And  light,  51,  ill. 

And  solution,  47. 

And  temperature,  113. 

Classes,  3. 

Chemical  attraction,  4. 
Chemical  calculations,  103. 
Chemical  change,  I,  2,  14,  47. 

And  ozone,  22. 
Chemical  compounds,  69. 
Chemical  energy,  in. 
Chemical  equivalents,  123. 
Chemicking,  138. 
Chemism,  4. 
Chemistry,  defined,  I,  2. 

Organic,  405. 
Chemists'  table,  447. 
Chest,  256. 

Chili  saltpeter,  231,  292. 
Chinese  white,  336. 
Chloral,  412. 

Hydrate,  412. 
Chloride,  of  lime,  137. 

Test,  144. 

Chlorides,  139,  140,  143,  144. 
Chlorination  process,  315. 
Chlorine,  133. 

Acids,  91. 

And  hydrogen,  136. 

And  water,  57. 

Available,  137. 

Compounds,  296. 

Determination  of  atomic  weight,  171 . 


Chlorine,  continued. 

Hydrate,  139. 

Liquid,  139. 

Name,  135. 

Nascent,  139. 

Occurrence,  133. 

Preparation,  133-134. 

Properties,  135. 

Uses,  139. 

Water,  51,  135. 
Chloroform,  412. 
Chlorophyll,  373. 
Chloroplatinic  acid,  394. 
Choke  damp,  203. 
Chroma.  365. 
Chromates,  366,  368. 
Chrome  alum,  350,  367,  368. 

Iron  ore,  365. 

Orange,  368. 

Red,  368. 

Steel,  366. 

Yellow,  367. 
Chromic  chloride,  368. 

Compounds,  368. 

Hydroxides,  368. 

Oxide,  368. 

Sulphate,  368. 
Chromite,  365. 
Chromites,  368. 
Chromium,  365. 

As  a  metal,  368. 

Compounds,  366,  368. 

In  minerals  and  rocks,  365. 

Name,  365. 

Ore,  365. 

Silicide,  258. 

Tests,  367,  368. 

Trioxide,  368. 

Uses,  365. 

Chromous  compounds,  368. 
Chronological  table  of  chemists,  447. 
Cinchona  tree,  433. 
Cinchonine,  433. 
Cinder,  375. 
Cinnabar,  338, 


564 


Index. 


Citric  acid,  419. 

Classification,  organic  compounds,  408. 

Periodic,  398. 
Clay,  351,  352. 

Aluminium  from,  345. 
Clouds,  67. 
Coal,  184. 

And  graphite,  118. 

Beds,  184,  185. 

Bituminous,  185,  186,  210. 

Composition,  186. 

Distillation,  210. 

Distribution,  187. 

Fields,  1 86,  187. 

Fire,  196. 

Gas,  210. 

Gas  plant,  211. 

Mines,  gases  in,  202,  203. 

Products  from,  213. 

Section,  185. 

Soft,  185,  189. 
Coal  tar,  213. 

Dyes,  431. 
Cobalt,  389. 

Blue,  390. 

Test,  390. 
Cocaine,  433. 
Coca  plant,  433. 
Coffee,  433. 
Coins,  gold,  316. 

Nickel,  306,  389. 

Silver,  311. 
Coke,  189,  213. 

As  fuel,  190. 

Coal  gas  from,  210. 

From  petroleum,  209. 

In  iron  smelting,  377. 
Colemanite,  262. 
Collodion,  428. 
Colored  glass,  260. 
Color  of  metals,  279. 
Combination,  3. 

By  volume,  53. 

By  weight,  53. 

Of  gases,  1 66. 


Combustion,  15,  1 6,  28,  67,  148,  191. 

And  flame,  217. 

Old  theory,  15. 

Products,  218. 

Spontaneous,  14. 

Common  salt,  133,  286,  287.     See  So- 
dium chloride. 
Complete  fertilizer,  271. 
Components,  8. 
Composition,  ammonia  gas,  153. 

Coal,  1 86. 

Carbon  dioxide,  195. 

Earth's  crust,  6. 

Heavenly  bodies,  404. 

Hydrochloric  acid,  143. 

In  per  cent,  103. 

Natural  waters,  38. 

Nitric  acid,  157. 

Nitric  oxide,  162. 

Nitrous  oxide,  161. 

Of  a  compound,  50. 

Organic  compounds,  405. 

Water,  2^  27,  57. 
Compounds,  chemical,  8,  69. 

Saturated,  177. 

Unsaturated,  177,  178. 
Concentrated,  defined,  41. 

Alum,  349. 

Concentration  of  ore,  281. 
Condenser,  39,  40. 

Coal  gas,  212. 

Iodine,  231. 
Conductivity,  metals,  279. 

Solutions,  126-127. 
Cones,  Bunsen  flame,  220. 

Flame,  216-217. 
Conservation,  energy,  in. 

Matter,  4,  104. 
Constitution,  benzene,  430. 
.    Organic  compounds,  406. 

See  Composition. 
Constitutional  formula,  407. 
Contact  method  for  sulphuric  acid,  24^ 
Converter,  381. 
Cooking  soda,  290. 


Index. 


565 


Copper,  301. 

Acetate,  417. 

Alloys,  305,  316. 

And  sulphuric  acid,  243. 

Arsenite,  243. 

Carbonates,  301,  302,  303. 

Coins,  306. 

Compounds,  306. 

Electrolytic,  303. 

Fluoride,  227. 

From  Michigan,  301,  302. 

Glance,  301. 

History,  301. 

Iron  sulphides,  301,  302,  373. 

Metallurgy,  302. 

Name,  301. 

Native,  301,  302. 

Nitrate,  307. 

Ores,  301. 

Oxides,  301,  302,  303,  306,  307. 

Production,  302. 

Properties,  303. 

Purification,  303. 

Pyrites,  301. 

See  Cupric  and  Cuprous. 

Smelting,  302. 

Region,  map,  374. 

Replacement,  304. 

Replacing  power,  304. 

Silicide,  258. 

Sulphate,  307. 

Sulphide,  301,  307. 

Test,  304. 

Uses,  304. 
Copperas,  385. 
Coquina,  322. 
Coral,  323. 
Cordials,  411. 
Corrosive  sublimate,  340. 
Corundum,  343,  346,  347. 
Cottolene,  421. 
Courtois,  231,  444. 
Crayon,  323. 
Cream  of  tartar,  290,  418. 

Potassium  carbonate  from,  297. 


Crockery,  352. 

Crocoisite,  365. 

Crocoite,  365. 

Crocus,  384. 

Crucible  process  for  steel,  380. 

Cruikshank,  119. 

Cryolite,  225,  343,  344,  350. 

Crystal,  rock,  255. 

Crystals,  44,  440. 

Hexagonal,  441. 

Isometric,  441. 

Monoclinic,  442. 

Orthorhombic,  239,  441. 

Production,  440. 

Snow,  35. 

Systems,  440. 

Tetragonal,  441. 

Triclinic,  442. 
Crystallization,  44,  440. 

Water  of,  45,  46. 
Cubic  cleavage,  363. 
Cupel,  310. 
Cupellation,  310,  360. 
Cupric  compounds,  306. 

Oxide,  307. 

Sulphate,  307. 

Sulphide,  307. 
Cuprite,  301,  306. 
Cuprium  aes,  301. 
Cuprous  compounds,  306. 

Oxide,  306,  426. 

Sulphide,  307. 
Cuprum,  301. 
Current,  electric,  121. 
Cyanic  acid,  198. 
Cyanide,  mercury,  198. 

Potassium,  198. 

Process,  198,  315. 

Iron,  387. 
Cyanogen,  198. 
Cymogene,  208. 
Cyprian  brass,  301. 

Dalton,  77,  79,  444. 

Davy,    53,    97,    119,    135,    161,    182, 


566 


Index. 


221,  231,  284,  286,   319,   343, 

444- 

Deacon's  process  for  chlorine,  134. 
Decay,  17,  67,  155,  191. 
Decomposition,  3. 

Double,  3,  45. 

Heat  of,  1 1 3. 

Definite  proportions,  law,  76. 
Deflagration,  159. 
Dehydrated,  46. 
Deliquescence,  46,  67. 
Destructive  distillation,  188,  202. 
Determination,  atomic  weights,  170- 

171. 

Developer,  313. 
Deville,  343. 
Dewar,  29,  444. 

Bulb,  70. 
Dew  point,  66. 
Dextrin,  427. 
Dextrose,  425,  426. 
Diacid  base,  94. 
Diamond,  181,  182,  190. 

Artificial,  182. 

Cheap,  257. 

Drill,  182. 

Diatomaceous  earth,  256,  257. 
Diatoms,  256. 
Dibasic  acid,  92. 
Dicalcium  phosphate,  271. 
Dichlorethane,  204. 
Dichromates,  368. 
Diffusion,  26,  68. 
Diluents,  215. 
Dilute,  41. 

Diphosphates,  269,  271. 
Disinfectant,  carbolic  acid,  431. 

Formaldehyde,  413. 
Disodium  phosphate,  269. 
Displacement,  downward,  135. 

Upward,  148. 
Dissociation,  by  heat,  151. 

Electrolytic,  125,  127. 
Distillate,  40. 
Distillation,  39. 


Distillation,  continued. 

Coal,  210. 

Destructive,  188,  202,  204. 

Dry,  147. 

Petroleum,  208. 

Water,  39. 

Wood,  1 88,  409. 
Distilled  liquors,  411. 

Water,  40. 
Dolomite,  331,  334. 
Double  decomposition,  3,  45. 

Refraction,  320. 
Downward  displacement,  135. 
Drinking  water,  39. 

Lead  in,  359. 
Drummond  light,  29. 
Ductile  metals,  279. 
Dulong,  444. 

And  Petit,  172. 
Dumas,  182,  397,  444. 

And  Boussingault,  66. 

And  Stas,  56,  57. 
Dutch  leaf,  305. 

Metal,  305. 

Process  for  white  lead,  361. 
Dyads,  176. 
Dyeing,  350. 
Dynamite,  422. 

Earthenware,  352. 

Effervescence,  42,  193. 

Effervescing  powder,  290. 

Efflorescence,  46. 

Electrical  conductivity,  126^ 

Electric  battery,  120. 

Electric  furnace,  114-115,  184,  365. 

Industrial  use  of,  116. 
Electricity  and  chemical  action,  119. 
Electric  light  carbons,  209. 
Electrochemical  equivalent,  123. 

Terms,  120. 
Electrochemistry,  119. 
Electrodes,  118,  120,  121,  184,  190. 
Electrolysis,  120. 

Aluminium  oxide,  343. 


Index. 


56? 


Electrolysis,  continued. 

And  solution,  126. 

Calcium  iodide,  319. 

Carnallite,  332. 

Copper  sulphite,  303. 

Galena,  358. 

Gold  solution,  315. 

Hydroxides,  284. 

Illustrations,  122. 

Industrial  application,  124. 

Metals,  281. 

Potassium  hydroxide,  294. 

Sodium  chloride,  122,  291. 

Sodium  hydroxide,  284. 

Sodium  nitrate,  302. 

Theory  of,  125. 

Water,  52,  123. 

Zinc  chloride,  122. 
Electrolytic  cell,  122. 

Copper,  303. 

Dissociation,  125. 

Process  for  chlorine,  134. 

Process  for  white  lead,  362. 

Separation  of  gold  and  silver,  315. 
Electro-negative  ions,  121,  122. 

Positive  ions,  121,  122. 

Silicon,  257. 

Thermal    manufacture    of    carbon 

disulphide,  252. 
Electrolyte,  120,  128. 
Electroplating,  1 24- 125. 
Electrotyping,  124-125. 
Elements,  5,  6,  7,  448,  449. 

Acid  properties,  396. 

Basic  properties,  396. 

Bivalent,  176. 

Classification,  396. 

Families,  397. 

General  relations,  396. 

In  earth's  crust,  6. 

In  organic  compounds,  405. 

In  sun,  404. 

Numerical  relations,  397. 

Periodic  classification,  398. 

Prediction,  401. 


Elements,  continued. 

Quadrivalent,  176. 

Quinquivalent,  176. 

Spectra,  402. 

Trivalent,  176. 

Table,  448,  449. 

Univalent,  176. 
Emerald,  347. 
Emery,  343,  346. 
Empirical  formula,  178,  407. 
Emulsin,  432. 
Endothermic,  112. 
Energy,  chemical,  4. 

Mechanical,  33. 
Enriching  gas,  213,  215. 
Epsom  salts,  333. 
Equation,  83,  84. 

Gas,  175. 

Illustrating  reactions,  106. 

Ionic,  129,  130. 

Molecular,  175. 

Problems  based  on,  107. 

Quantitative  significance,  104. 

Thermal,  112. 

Volumetric,  175. 
Equivalents,  100. 

And  valence,  178. 

Chemical,  123. 

Electrochemical,  123. 

Multiples,  101. 

Table,  100. 
Erosion,  32. 
Esters,  419. 
Etching,  227. 
Ethane,  202,  203,  409. 

Graphic  formula,  407. 
Ether,  ethyl,  413. 

And  water,  43. 

Sulphuric,  414. 
Ethereal  salts,  419,  420. 
Ethers,  413. 
Ethyl,  406,  409. 

Acetate,  419. 

Alcohol,  406,  408,  409. 

Butyrate,  419. 


568 


Index. 


Ethyl,  continued. 
Ether,  413. 
Oxide,  414. 
Sulphuric  acid,  414. 
Ethylene,  204. 
Chloride,  204. 
In  illuminating  gas,  216. 
Series,  202. 
Eudiometer,  53,  54. 
Evaporation,  440. 

Exercises,  9,  20,  30,  48,  58,  73,  85,  98, 
108,   130,   145,   163,  178,  198, 
222,  233,  253,  263,  276,  282, 
299,   3*7»  329,  340,  352»  363» 
372,  390,  394,  404,  433- 
Exhauster,  212. 
Exothermic,  112. 
Exposure,  photographic,  313. 

Factors,  83. 

Fahrenheit  thermometer,  439. 

Families  of  elements,  397. 

Faraday,  120,  123,  128,  139,  193,  444. 

Law,  123. 
Fats,  420. 

Fatty  acid  series,  414. 
Fehling's  solution,  426. 
Feldspar,  293,  343,  351. 
Fermentation,  192,  410. 

Acetic,  416. 

Alcoholic,  410. 

In  bread  making,  421. 

Sugar,  410. 
Ferments,  410. 

And  glucosides,  432. 
Ferric  compounds,  384. 

Chloride,  386. 

Ferrocyanide,  388. 

Hydroxides,  385. 

Oxide,  384. 

Sulphate,  385. 

Sulphide,  386. 
Ferricyanides,  387,  388. 
Ferrochrome,  366. 
Ferrocyanides,  387,  388. 


Ferro-ferric  oxide,  385. 
Ferromanganese,  369,  378. 
Ferrous  compounds,  384. 

Carbonate,  387. 

Chloride,  386. 

Ferric  oxide,  384. 

Ferricyanide,  388. 

Ferrocyanide,  388. 

Hydroxide,  385. 

Sulphate,  385. 

Sulphide,  240,  385. 
Ferrum,  373. 
Fertilizer,  73,  271. 

Manufacture,  271. 

Potassium  salts  as,  298. 

Sodium  nitrate  as,  292. 
Film,  photographic,  313. 
Filter,  charcoal,  188. 
Filtering  water,  39. 
Fire,  15. 

Damp,  202. 

Extinguisher,  194. 
Fireworks,  14,  332. 
Fixed  air,  196. 

Alkalies,  93. 

Fixing,  in  photography,  303. 
Flame,  216. 

Acetylene,  207,  216. 

And  combustion,  217. 

Bunsen,  219. 

Hydrogen,  27,  112. 

Non -luminous,  219,  220. 

Oxidizing,  221,  222. 

Oxyhydrogen,  29. 

Parts,  216-217. 

Reducing,  222. 

Smoky,  218. 
Flashing  point,  209. 
Flavors,  419. 
Flint,  255,  256. 
Flour,  wheat,  427. 
Flowerpots,  352. 
Flowers  of  sulphur,  238. 
Fluid,  magnesia,  334. 
Fluorides,  227,  343. 


Index. 


569 


Fluorine,  225. 

Apparatus,  226. 

Isolation,  225. 

Liquid,  226. 

Name,  225. 

Properties,  226. 
Fluor  spar,  225,  226. 
Flux,  281,  375. 
Food,  water  in,  31,  32. 
Fool's  gold,  386. 
Formaldehyde,  412. 
Formalin,  413. 
Formation,  heat  of,  1 12. 
Formula,  82. 

Constitutional,  407. 

Empirical,  178,407. 

Graphic,  178,  407,  413,  414. 

Molecular,  174. 

Rational,  407. 

Simplest,  104,  174,  175. 

Structural,  178,  407,  413,  414. 
Fossil,  from  coal  bed,  185. 
Frame,  for  soap,  423. 
Franklinite,  334. 

French  process  for  white  lead,  362. 
Fructose,  425. 
Fruit  sugar,  425. 
Fuming  acid,  nitric,  163. 

Sulphuric,  251. 
Furnace,  blast,  281,  375. 

Reverberatory,  281,  282. 
Fusible  alloys,  337,  360. 

Metals,  275. 
Fusion,  for  crystals,  440. 

Gahnite,  334. 
Galena,  357,  362. 
Crystals,  362. 
Gallic  acid,  432. 
Gallium,  401. 
Galvanic  cell,  119. 
Galvanized  iron,  336. 
Gangue,  280. 

Gaps  in  periodic  system,  401. 
Garnet,  347. 


Gas,  61. 

Carbon,  190,  213. 

Coal,  210. 

Effect  of  heat  on  volume,  18,  19. 

Effect  of  pressure  on  volume,  18. 

Equation,  175. 

Flame,  structure,  218. 

Holder,  212. 

Illuminating,  210. 

Marsh,  202. 

Natural,  209. 

Producer,  25. 

Sylvestre,  196. 

Volume,  reduction,  53,  54. 

Water,  25,  196,  213. 

Water,  plant,  214. 
Gases,  absorption  by  charcoal,  188. 

By  platinum,  394. 

Combination  by  volume,  166. 

Inert,  69. 

In  mines,  221. 

Properties,  166. 

Solution  of,  41. 
Gasolene,  208. 
Gay-Lussac,  55,  231. 

Law,  1 66. 

Tower,  249. 

Gelatine  plate  and  film,  313. 
Gems,  aluminium,  347. 

Artificial,  347. 

Glass,  260. 

Quartz,  257. 
Generator,  acetylene,  207. 

Water  gas,  213. 
German  process  for  white  lead,  362. 

Silver,  306,  389. 
Geyserite,  258. 
Gin,  411. 
Glacial  acid,  acetic,  415. 

Phosphoric,  268. 
Glass,  258. 

And  hydrofluoric  acid,  227. 

Annealing,  260. 

Blasting,  257. 

Blowing,  259,  260. 


570 


Index. 


Glass, 

Bohemian,  260} 

Colored,  260. 

Constituents,  259. 

Crown,  260. 

Cut,  260. 

Flint,  260. 

Kinds,  258,  259. 

Manufacture,  259. 

Plate,  259. 

Polishing,  260. 

Production,  260. 

Typical  mixture,  259. 

Window,  259. 
Glauber,  140,  445. 

Salt,  292. 

Glazing  pottery,  352. 
Globigerina  ooze,  322,  323. 
Glover  tower,  248. 
Glucose,  425,  426. 
Glucosides,  432. 
Glycerides,  420. 
Glycerine,  420. 

Preparation,  422. 

Properties,  421. 

Relation  to  soap,  420. 

Uses,  421. 
Glycerol,  422. 
Glyceryl,  429. 

Oleate,  420. 

Palmitate,  420. 

Stearate,  420. 
Gneiss,  255. 

Gogebic  iron  range,  374. 
Gold,  314. 

Alloys,  314,  316. 

Amalgam,  339. 

Chloride,  315,  316,  317. 

Coin,  316. 

Compounds,  317. 

Cyanide,  315,  317. 

Distribution,  309,  314. 

Dust,  314. 

Dutch,  305. 

Finely  divided,  317. 


Gold,  continued. 

Fool's,  386. 

History,  313. 

Leaf,  316. 

Making,  314. 

Map  of  distribution,  309. 

Name,  314. 

Nugget,  314. 

Parting,  315. 

Pen  tips,  394. 

Plating,  317. 

Production,  314. 

Properties,  316. 

Purification,  315. 

Red,  316. 

Reduction  of  compounds,  317. 

Separation  from  silver,  315. 

Test,  317. 

Uses,  316. 

White,  316. 
Graham,  26,  445. 
Gram,  437. 
Granite,  255. 
Grape  sugar,  425,  426. 
Graphic     formula,     178,      407,     413, 

414. 
Graphite,  183,  190. 

Artificial,  118. 
Gravimetric,  53. 

Composition,  air,  66. 

Composition,  water,  55,  57. 
Gray  cast  iron,  378. 
Green  fire,  329. 

Pigments,  368. 

Vitriol,  246,  385. 
Grindstones,  256. 
Groups  of  elements,  397. 
Guano,  271,  331. 
Guignet's  green,  368. 
Gun  cotton,  428. 

Metal,  305. 
Gunpowder,  14,  296. 

Smokeless,  428. 
Gypsum,  326. 

Reduction  of,  235. 


Index. 


Haemoglobin,  373. 
Halides,  96,  225. 
Hall,  343- 

Process  for  aluminium,  343,  344. 
Halogens,  225. 
Haloid  salts,  225. 
Hardness,  of  metals,  279. 

Of  water,  37,  327. 

Permanent,  327. 

Temporary,  327. 
Hard  water,  37. 

Coal,  185,  1 86. 
Harveyized  steel,  380. 
Hausmannite,  369. 
Heat,     and     chemical     action,     112, 

H3- 

And  oxidation,  14. 

From  burning  hydrogen,  1 12. 

In  electric  furnace,  114. 

Of  decomposition,  113. 

Of  formation,  112. 

Of  neutralization,  130. 
Heavenly  bodies,  constitution,  404. 
Helium,  69,  404. 
Hematite,  373. 
Henry's  law,  42. 

Heroult  process  for  aluminium,  344. 
Hexagonal  crystals,  441. 
Hofmann,  445. 

Apparatus,  52. 
Honey,  425. 
Horn  silver,  133,  308. 
Humboldt,  55. 
Hydrargyrum,  338. 
Hydrate,  93. 

Chlorine,  139. 
Hydrated,  46. 
Hydraulic  lime,  325. 

Main,  210. 

Mining,  314. 
Hydriodic  acid,  232. 
Hydrocarbons,  202,  408. 
Hydrobromic  acid,  230. 
Hydrochloric  acid,  140-143, 

Commercial,  141. 


Hydrochloric,  continued. 

Composition,  143. 

Liquefied,  142. 

Test,  144. 

Hydrocyanic  acid,  198. 
Hydrofluoric  acid,  227,  257. 

Vapor  density,  228. 
Hydrogen,  23. 

And  chlorine,  136. 

And  periodic  classification,  401. 

And  steam,  24. 

And  water,  50. 

Arsenide,  273. 

Chemical  conduct,  27. 

Diffusion,  26. 

Dioxide,  59. 

Discovery,  30. 

Explosions,  28. 

Flame,  27. 

In  acids,  24,  87,  90. 

Ions,  121. 

Liquid,  29. 

Name,  25,  30. 

Peroxide,  59. 

Physical  properties,  25. 

Preparation,  24. 

Solid,  29. 

Valence,  176. 

Weight  of  liter,  25.  + 
Hydrogen  sulphide,  240,  241,  242. 

Composition,  241.  .-- 

Test,  242. 

Water,  241. 
Hydroquinone,  431. 
Hydroxides,  89,  93. 

And  alcohols,  409. 

Common  names,  93. 

Organic,  409. 
Hydroxyl,  89,  94. 
Hygroscopic,  46. 
Hypo,  91,  252. 
Hypophosphites,  269. 
Hyposulphite  in  photography,  313. 
Hypothesis,  76. 

Avogadro's,  167. 


572 


Index. 


Ice,  32,  34,  35. 

Making  plant,  153. 

Manufactured,  153. 

Stone,  350. 
Iceland  spar,  320. 
Illuminants,  216. 
Illuminating  gas,  210. 

Carbon  monoxide  in,  197. 

Characteristics,  215. 

Composition,  215. 

Illuminating  power,  216. 

Impurities,  240. 

Luminosity,  216. 
Indicator,  98. 

Inert  gases  in  atmosphere,  69. 
Infusorial  earth,  256,  257. 
Ingots,  381. 
Ink,  385,  418. 

Indelible,  312. 

Printer's,  190. 

Writing,  433. 

Inorganic  compounds,  405. 
Insoluble  substances,  41. 

Sulphate,  test,  251. 
Intervals    in    periodic    classification, 

398. 

Iodides,  232. 
Iodine,  230. 

Commercial  preparation,  231. 

Compounds,  232. 

Detection,  232. 

Determination,  252. 

Discovery,  231. 

In  seaweed,  230. 

Name,  232. 

Preparation,  230. 

Production,  233. 

Properties,  231. 

Purification,  231. 

Source,  293. 

Test,  232. 

Uses,  233. 

Vapor  density,  232. 
lodoform,  233,  412. 
Ionic  equation,  129,  130. 


lonization,  125. 

And  acids,  bases,  and  salts,  129. 

Application,  129. 

Table,  127. 
Ions,  120,  121,  125,  126. 

Test  for,  129. 

Iridium,  226,  392,  393,  394. 
Iridosmine,  394. 
Iron,  373. 

Acetate,  417. 

Alum,  386. 

And  coke,  190. 

By  alcohol,  383. 

By  hydrogen,  383. 

Carbide,  209,  285. 

Carbonate  ores,  374. 

Cast,  377,  378. 

Chemistry  of  smelting,  377. 

Chlorides,  386. 

Compounds,  384.     See  Ferric  and 
Ferrous. 

Cyanides,  387. 

Disulphide,  386. 

Galvanized,  386. 

History,  373. 

Impurities,  377. 

Liquor,  417. 

Magnetic  oxide,  385. 

Malleable,  879. 

Map  of  deposits,  374. 

Metallurgy,  379. 

Ore,  373,  374. 

Ore,  chrome,  365. 

Ore,  consumption,  377. 

Ore,  deposits,  374. 

Ore,  reduction,  197,  375. 

Oxides,  384. 

Passive,  384. 

Pig,  377- 
Properties,  383. 
Pyrites,  373,  385,  386. 
Rust,  383. 
Rusting,  14. 
Silicide,  258. 
Smelting,  375. 


Index. 


573 


Iron,  continued. 

Spiegel,  369. 

Sulphides,  386. 

Symbol,  373. 

Test,  388. 

Varieties,  377. 
Isomerism,  204. 
Isomers,  204. 
Isometric  crystals,  441. 
Ivory  black,  189. 

Jasper,  256. 
Javelle's  water,  139. 

Kainite,  294,  298,  331. 

As  fertilizer,  298. 
Kali,  204. 
Kalium,  294. 
Kaolin,  351,  352. 
Kassiteros,,  354. 
Kelp,  231. 
Kerosene,  209. 
Kieserite,  331,  333. 
Kilogram,  437. 

Kindling  temperature,  113,218,  221. 
Kirchhoff,  403,  445. 
Krypton,  69,  404. 

Labarraque's  solution,  139. 
Lactic  acid,  290,  418. 
Lactose,  425. 
Lake,  350. 
Lampblack,  190. 
Laudanum,  433. 

Lavoisier,  5,  15,  16,  18,  25,  27,  30,  50, 
55,  63,  64,  88,  97,   157,    182, 
196,  396,  445. 
Law,  75. 

Boyle,  19. 

Charles,  19. 

Conservation  of  energy,  in. 

Definite  proportions,  75,  76,  79. 

Faraday,  123. 

Gay-Lussac,  166. 

Henry,  42. 


Law,  continued. 

Matter,  5. 

Multiple  proportions,  77,  78. 

Periodic,  398. 

Specific  heat,  172. 
Lead,  357. 

Acetate,  363,417. 

Alloys,  360. 

Argentiferous,  308. 

Black,  183,  359. 

Carbonate,  357,  361. 

Carbonate,  basic,  361. 

Chambers,  249. 

Chloride,  363. 

Chromate,  367. 

Chromate,  native,  365. 

Compounds,  363. 

Compounds,  poisonous,  359. 

Cupellation  process,  310. 

Dioxide,  361. 

History,  357. 

Hydroxide,  362. 

In  drinking  water,  359. 

Interaction  with  metals,  359. 

Metallurgy,  358. 

Monoxide,  360. 

Nitrate,  363. 

Nitrate,     behavior      with      heat, 
163. 

Ore,  357. 

Oxides,  360. 

Parkes  process  for,  309. 

Pencils,  184. 

Peroxide,  361. 

Phosphate,  357. 

Pipe,  366. 

Production,  357. 

Properties,  358. 

Silver  bearing,  308. 

Spongy,  360. 

Sugar  of,  363,  417. 

Sulphate,  357,  363. 

Sulphide,  242,  357,  362,  363. 

Test,  363, 

Tetroxide,  360. 


574 


Index. 


Lead,  continued. 

Uses,  359. 

White,  361. 
Leather,  433. 
Leblanc  process  for  sodium  carbonate, 

288. 

Lemon  juice,  90. 
Levulose,  425. 
Liebig,  230,  445. 
Life  and  carbon  dioxide,  194. 

Oxygen,  16. 

Nitrogen,  72. 

Phosphorus,  270. 

Potassium,  298. 

Light     and      chemical     action,     51, 
ill. 

Silver  salts,  312,313. 
Lignite,  185. 
Lime,  324. 

Air  slaked,  324. 

And  water,  113,  324. 

Caustic,  324. 

Chloride  of,  137. 

Hydraulic,  325. 

Light,  29,  324. 

Making,  192,  324,  325. 

Milk  of,  326. 

Quick,  324. 

Superphosphate,  271. 

Uses,  324. 

See  Calcium  oxide. 
Limekiln,  193,  325. 
Limestone,  320. 

As  flux,  377. 

Burning,  325. 

Caves,  321,  322. 

Fossil,  322. 

Solubility,  321. 

Uses,  323. 
Lime  water,  325. 

And  carbon  dioxide,  192,  325. 

Detection,  68. 

Preparation,  326. 

See  Calcium  hydroxide. 
Liming,  138. 


Limonite,  373. 
Links,  407. 
Liquid  air,  12,  69. 

Acetylene,  205. 

Ammonia,  148-149,  153. 

Carbon  dioxide,  193. 

Chlorine,  139. 

Fluorine,  226. 

Hydrogen,  29. 

Oxygen,  18. 

Sulphur  dioxide,  244. 
Liquids,  solubility,  43. 
Liquor,  alcoholic,  411. 

Distilled,  411. 

Iron,  417. 

Red,  350. 

List  of  reference  books,  450. 
Litharge,  360. 
Lithia  water,  298. 
Lithium,  298. 

Citrate,  298. 

Discovery,  294. 

Test,  298. 
Litmus,  action  on,  acid,  90. 

Alkali,  92. 

Base,  92. 

Neutral  substance,  94. 

Salt,  94. 
Loadstone,  385. 
Lubricating  oil,  209. 
Luminosity,  illuminating  gas,  2 1 6. 

Of  flame,  218. 
Luminous  paint,  329. 
Lunar  caustic,  312. 
Luray  cavern,  321,  322. 
Luster,  279. 

Madder,  432. 
Magnalium,  346. 
Magnesia,  333,  334,  370. 

Alba,  334,  370. 

Black,  370. 

Fluid,  334. 

Mixture,  333. 

Nigra,  370. 


Index. 


575 


Magnesia,  continued. 

Stone,  370. 

Uses,  333. 
Magnesite,  334. 
Magnesium,  331. 

Alloy,  346. 

Bromide,  228. 

Calcium  carbonate,  331. 

Carbonate,  331,  334. 

Chloride,  333. 

Citrate,  334. 

Compounds  in  soil,  331. 

Compounds  and  water,  327. 

Hydroxide,  333. 

Nitride,  153,332. 

Oxide,  333.     See  Magnesia. 

Phosphates,  331. 

Preparation,  332. 

Properties,  332. 

Ribbon,  332. 

Sulphate,  333. 

Uses,  332. 

Magnetic  oxide  of  iron,  385. 
Magnetite,  373,  385. 
Majolica,  352. 
Malachite,  301,  308. 
Malic  acid,  418. 
Malleable  iron,  379. 

Metals,  279. 
Mammoth  cave,  322. 
Manganates,  371. 
Manganese,  369. 

Alloys,  369. 

As  non-metal,  371. 

Black  oxide,  370. 

Compounds,  371. 

Dioxide,  369. 

Isolation,  370. 

History,  370. 

Name,  370. 

Ores,  369. 

Preparation,  369. 

Production,  369. 

Properties,  369. 

Test,  372. 


Manganese,  continued. 

Uses,  369. 
Manganesium,  370. 
Manganite,  369. 
Manganous  compounds,  371. 

Chloride,  370,  371. 

Hydroxide,  370. 

Sulphate,  371. 

Sulphide,  371. 
Mantle,  Welsbach,  222. 
Map,  copper  deposits,  374. 

Gold,  309. 

Iron,  374. 

Silver,  309. 
Marble,  320. 
Marchand  tube,  56. 
Marengo  cave,  322. 
Marquette  iron  range,  374. 
Marsh  gas,  202. 
Marsh's  test  for  arsenic,  273. 
Massicot,  360. 
Matches,  270. 
Matte,  copper,  302. 
Matter,  conservation,  4. 

Properties,  i,  2. 
Meadowsweet,  430. 
Meerschaum,  331. 
Mendeleeff,  398,  445. 
Menominee  iron  range,  374. 
Mercuric  chloride,  340,  357. 

Cyanide,  198. 

Nitrate,  340. 

Oxide,  1 8,  339. 

Sulphide,  338,  340. 
Mercurous  chloride,  339,  357. 

Nitrate,  340. 
Mercury,  337. 

Alloys,  339. 

Compounds,  339. 

Deposits,  338. 

Fulminating,  339. 

Name,  338. 

Native,  337. 

Ore,  338. 

Preparation,  338, 


576 


Index. 


Mercury,  contintced. 

Production,  338. 

Properties,  338. 

Purification,  338. 

Specific  heat,  172. 

Transportation,  338. 

Uses,  339. 

Vapor  density,  169,  339. 
Mesabi  iron  range,  374. 
Metal,  and  non-metal,  278. 

Babbit's,  336. 

Bath,  305. 

Bell,  306. 

Britannia,  306,  356. 

Dutch,  305. 

Gun,  305. 

Hypothetical,  150. 

Muntz,  305. 

Newton's,  275. 

Rose's,  275. 

Speculum,  306. 

Type,  360. 

White,  306. 

Wood's,  275,  337. 
Metallic  ions,  121. 

Luster,  279. 
Metalloids,  278. 
Metallurgy,  280. 

Copper,  302. 

Lead,  358. 

Iron,  375. 

Silver,  309,  310. 
Metals,  action  with  nitric  acid,  158. 

Alkali,  284. 

Alkaline  earth,  319. 

Antifriction,  336. 

Chemical  properties,  279. 

Classification,  396. 

Familiar,  7. 

Found  free,  280. 

General  properties,  278. 

Known  to  ancients,  280. 

Occurrence,  279. 

Physical  properties,  278. 

Platinum,  394. 


Metals,  continued. 

Preliminary  treatment,  280. 

Preparation,  280. 
Metamerism,  204. 
Metaphosphates,  269. 
Metaphosphoric  acid,  268. 
Metastannic  acid,  355. 
Metathesis,  3. 
Meter,  defined,  437. 

Gas,  212. 
Methane,  202,  409. 

Graphic  formula,  407. 

In  natural  gas,  209. 

Series,  202. 
Methyl,  406,  409. 

Alcohol,  409. 

Benzene,  430. 

Salicylate,  432. 
Methylated  spirit,  410. 
Metric  abbreviations,  438. 

Apparatus,  394. 

Equivalents,  438. 

System,  437. 

Ton,  32. 

Transformations,  438. 
Mexican  onyx,  322. 
Meyer,  Lothar,  398,  445. 
Mica,  293,  343. 
Microcosmic  salt,  269. 
Milk  of  lime,  326. 

Sulphur,  240. 

Milner's  process  for  white  lead,  362. 
Mineral,  defined,  280. 

Compounds,  405. 

Springs,  37,  42. 

Water,  37. 
Minerals,  258. 
Minium,  360. 
Mispickel,  272. 
Mixture,  9,  77. 

Air,  69. 
Modification,  allotropic,  191. 
Moissan,  114,  116,  182,  184,  225,  226, 

3!9.  365,  445- 
Moissan's  electric  furnace,  1 14. 


Index. 


577 


Molecular  equation,  175. 

Formula,  174. 
Molecular  weights,  103,  128,  168. 

And  vapor  density,  168. 

Determination,  170,  171. 

Exact,  170. 

Hydrogen,  169. 

Steam,  169. 
Molecules,  80-8 1,  167-168. 

And  atoms,  80. 

And  equations,  175. 
Molybdenum,  369. 
Monacid  base,  94. 
Monads,  176. 
Monobasic  acids,  92. 
Monocalcium  phosphate,  271. 
Monoclinic  crystals,  442. 

Sulphur,  239. 
Monophosphates,  269. 
Mordants,  350,  357,  367. 
Morphine,  433. 
Mortar,  326. 
Moth  balls,  432. 
Mother  liquor,  230,  231. 
Mucilage,  427. 
Multiple  proportions,  law,  77-78. 

Table,  78. 
Muntz  metal,  305. 
Muria,  140. 

Muriate  of  ammonia,  151. 
Muriatic  acid,  92,  140. 
Muscovado  sugar,  424. 
Mutton  fat,  420. 

Naphtha,  208. 
Naphthalene,  432. 
Nascent  state,  138. 
Natrium,  284. 
Natron,  284. 
Natural  gas,  209. 
Natural  groups,  400. 

Waters,  38. 
Nature  of  solution,  48. 
Negative  electrode,  121. 

Photographic,  313. 


Neon,  69,  404. 
Neutral,  94. 

Reaction,  94. 
Neutralization,  88,  89,  97. 

And  ionic  theory,  130. 

Heat  of,  130. 
Newton's  metal,  275. 
Niagara  Falls,  industries  at,  1 1 6,  117, 

118,  155,  291,  344. 
Nicholson  and  Carlisle,  53,  119. 
Nickel,  388. 

Alloys,  306. 

Carbonyl,  198. 

Coin,  306,  389. 

Hydroxide,  389. 

Ores,  388. 

Plating,  389. 

Properties,  389. 

Steel,  383,  389. 

Test,  389. 

Uses,  389. 
Nickeloid,  389. 
Nicotine,  433. 
Niter,  72. 

Meal,  295. 

Source,  295. 
Nitrates,  158. 

Behavior  with  heat,  159. 

Deposits,  155. 

Test,  159. 
Nitric  acid,  154,  155,  156. 

Action  with  metals,  158. 

And  copper,  159,  162. 

And  electric  sparks,  155. 

Composition,  157. 

Formation,  155. 

Fuming,  163. 

Preparation,  155. 

Test,  159. 

Uses,  157. 
Nitric  oxide,  159,  162. 

Composition,  162. 
Nitrides,  72. 

Magnesium,  153. 
Nitrification,  155. 


578 


Index. 


Nitrites,  159. 
Nitrogen,  72. 

Discovery,  63. 

Effect  on  flame,  220. 

In  atmosphere,  63. 

Name,  72. 

Oxides,  78,  1 60. 

Pentoxide,  163. 

Peroxide,  159,  162,  163. 

Preparation,  72. 

Properties,  63,  72. 

Proportion  in  air,  64. 

Relation  to  life,  72. 

Tetroxide,  163. 

Trioxide,  163. 

Valence,  177,  178. 
Nitrous  acid,  159. 
Nitrous  oxide,  160,  161. 

Composition,  161. 

Discovery,  161. 
Nitrobenzene,  430. 
Nitroglycerine,  422. 
Nitrosyl-sulphuric  acid,  248. 
Nomenclature,  acids,  91. 

Bases,  93. 

Hydroxides,  93. 

Salts,  95. 

Non-luminous  flame,  219,  220. 
Non-metallic  ions,  121. 
Non-metals,  88. 

Classification,  396. 

General  properties,  278. 
Nordhausen  sulphuric  acid,  252. 
Normal  bismuth  nitrate,  276. 
Normal  salts,  96. 
Nugget,  gold,  314. 

Occlusion,  26. 
Ocean  water,  38. 

Salts  in,  38. 
Oil,  and  water,  43. 

Lamp  flame,  218. 

Lubricating,  209. 

Of  bitter  almonds,  431,  432. 

Of  vitriol,  92,  246. 


Oils,  420. 

Olefiant  gas,  204. 

Olein,  420. 

Oleomargarine,  421. 

Olive  oil,  420,  421. 

Onyx,  255. 

Opal,  256. 

Opaque,  279. 

Open  hearth  process  for  steel,  382. 

Opium,  433. 

Orange  mineral,  361. 

Ore,  defined,  280. 

Calcination,  281. 

Classes,  280. 

Dressing,  281. 
Organic  acids,  92,  414. 

Chemistry,  405. 

Compounds,  405,  406,  408. 
Orpiment,  272,  273. 
Orthophosphoric  acid,  268. 
Orthorhombic  crystals,  441. 

Sulphur,  239. 
Osmium,  392,  394. 
Ostwald,  445. 
Oxalic  acid,  417. 
Oxidation,  14,  192,  357. 

And  decay,  17. 

By  potassium  permanganate,  371, 

Of  food,  1 6. 
Oxide,  carbonic,  197. 
Oxides,  15. 

Acidic,  97. 

Basic,  97. 

Of  nitrogen,  160,  246-248. 

Relation  to  acids  and  bases,  96. 
Oxidized  silver,  311. 
Oxidizing  agent,  14,  60. 

In  matches,  270. 
Oxidizing  flame,  221,  222. 
Oxychloride,  antimony,  275. 

Bismuth,  276. 
Oxygen,  11. 

Absorption  by  silver,  311. 

And  blood,  16,  17. 

And  combustion,  15. 


Index. 


579 


Oxygen,  continued. 

And  flames,  218. 

And  ozone,  22. 

And  water,  51. 

Breathing  pure,  17. 

Erin's  process,  12. 

Discovery,  18. 

In  acids,  87,  88,  91,  97. 

In  atmosphere,  63. 

Liquid,  1 8. 

Name,  18,  88. 

Nascent,  138. 

Preparation,  II,  293. 

Properties,  12. 

Relation  to  life,  16. 

Solid,  1 8. 

Uses,  17. 

Weight  of  liter,  18. 
Oxyhydrogen  blowpipe,  17,  28,  29. 
Oxymuriate,  tin,  357. 
Ozone,  21,  113. 

In  atmosphere,  62. 

Formula,  169. 

Paint,  black,  190. 

Blue,  417. 

Lead,  357. 

Luminous,  329. 

Red,  273,  340,  361,  384. 

White,  336,  362. 

Yellow,  273,  367. 
Pakfong,  306. 
Paktong,  306. 
Palladium,  392,  394. 

Absorption  by,  26,  394. 
Palmitic  acid,  417. 
Palmitin,  420. 
Palm  oil,  417. 
Paper,  making,  429. 

Parchment,  428. 
Paracelsus,  30. 
Paraffin,  series,  203. 

Wax,  209. 
Paregoric,  433. 
Pads  green,  273,  417. 


Parkes  process  for  silver,  309. 
Parting,  gold  and  silver,  315. 
Passive  iron,  384. 
Paste,  gems,  262. 

Glass,  260. 

Starch,  427. 
Pastry,  raising,  290. 
Pearlash,  297. 
Peat,  185. 
Pentads,  176. 

Percentage  composition,  103. 
Periodic  classification,  398. 

Gaps,  401. 
Periodic  law,  398. 
Periodic  process  for  bromine,  229. 
Periodic  table  of  elements,  399. 
Periods  in  periodic  classification,  398. 
Permanent  hardness,  327. 
Peroxide,  hydrogen,  59. 

Sodium,  293. 
Petit,  172,445. 
Petrified  wood,  256,  257,  258. 
Petroleum,  207-209. 

Origin,  209. 

Production,  209. 

Refining,  208. 
Pewter,  356,  360. 
Phenol,  431. 

Derivatives,  431. 
Phenyl,  406. 

Methane,  406. 
Philosopher's  stone,  314. 
Phlogiston,  15,  1 8. 
Phosgene,  198. 
Phosphates,  265,  269. 

Acid,  26$. 

Dicalcium,  271. 

Disodium,  269. 

Monocalcium,  271. 

Primary,  269. 

Rock,  271. 

Secondary,  269. 

Slag,  271. 

Tricalcium,  271. 
Phosphine,  269. 


58° 


Index. 


Phosphonium  compounds,  269. 
Phosphor  bronze,  305. 
Phosphoric  acids,  268. 

Oxide,  268. 
Phosphorite,  265. 
Phosphorous  oxide,  268. 
Phosphorus,  265. 

Acids,  268. 

Action  on  air,  65,  72. 

And  ozone,  21. 

And  plants,  270. 

Black,  267. 

Discovery,  265. 

Electrolytic  manufacture,  266. 

In  plants  and  animals,  265. 

Manufacture,  265,  266. 

Minor  compounds,  269. 

Name,  267. 

Ordinary,  266. 

Oxides,  268. 

Pentachloride,  270. 

Pentoxide,  65,  268. 

Properties,  266. 

Purification,  266. 

Red,  267. 

Relation  to  life,  270. 

Salts,  268. 

Trichloride,  270. 

Uses,  267. 

Vapor  density,  169,  267. 

Yellow,  266. 
Photography,  ill,  312. 
Photometer,  216. 
Phylloxera,  240. 
Physical  changes,  I,  2. 
Pickles,  90,  416. 
Picrates,  431. 
Picric  acid,  431. 
Picromerite,  294. 
Pig  iron,  377. 
Pinchbeck,  305. 
Placer  mining,  314. 
Plants  and  atmosphere,  194. 

And  nitrogen,  72,  73. 

And  phosphorus,  270. 


Plants  and  atmosphere,  continued. 

And  potassium,  298. 

And  silica,  257. 
Plaster,  326. 

Of  Paris,  327. 
Plata,  392. 
Plate,  developing,  313. 

Photographic,  312. 
Platina,  392. 
Platinic  chloride,  394. 
Platinum,  392. 

Absorption  of  gases,  394. 

Alloys,  394. 

And  aqua  regia,  392. 

And  iridium,  226,  392. 

And  sulphur  dioxide,  245. 

And  sulphuric  acid,  249. 

Arsenide,  392. 

Black,  394. 

Compounds,  394. 

Discovery,  392. 

Dish,  393. 

Foil,  393. 

In  electric  light  bulbs,  393. 

Metals,  394,  401. 

Name,  392. 

Native,  392. 

Ore,  392. 

Preparation,  392. 

Print,  394. 

Production,  392. 

Properties,  393. 

Sheet,  393. 

Source,  392. 

Spongy,  392,  393. 

Uses,  393. 
Plumbago,  183. 
Plumbum,  357. 

Nigrum,  357. 
Polyhalite,  294. 
Polymerism,  206. 
Polymers,  206. 
Porcelain,  352. 
Portland  cement,  325. 
Positive  electrode,  121. 


Index. 


Potash,  297. 

Name,  294. 

Red  prussiate,  387. 

Yellow  prussiate,  387. 
Potassium,  293. 

Alum,  349. 

Antimonyl  tartrate,  274. 

Bichromate,  366. 

Bromide,  230. 

Carbonate,  297. 

Chlorate,  u,  12,  296,  297. 

Chloride,  295. 

Chloroplatinate,  394. 

Chromate,  366. 

Chromium  sulphate,  368. 

Cyanide,  198,  298,  315. 

Dichromate,  366. 

Discovery,  284. 

Ferricyanide,  387. 

Ferrocyanide,  198,  387. 

Hydroxide,  297,  298. 

Hypochlorite,  139. 

Iodide,  232,  233. 

Manganate,  371. 

Name,  294. 

Nitrate,  155,  295. 

Nitrite,  295. 

Permanganate,  370. 

Preparation,  294. 

Preservation,  294. 

Properties,  294. 

Relation  to  life,  298. 

Salts  and  starch,  298. 

Salts  at  Stassfurt,  293. 

Silicate,  258. 

Sulphate,  298. 

Sulphocyanate,  198, 

Tartrate,  418. 

Test,  294. 
Pottery,  352. 
Powder,  gun,  296. 

Smokeless,  428. 
Precipitate,  45. 
Precipitation,  44. 
Prefix,  centi-,  437. 


Prefix,  continued. 

Deca-,  437. 

Deci-,  437. 

Hecto-,  437. 

Hydro-,  91,  95. 

Hypo-,  91. 

Kilo-,  437. 

Milli-,  437. 

Per-,  91,  95. 
Press  cake,  296. 
Pressure,  normal,  18,  19. 
Priestley,  n,  16,  18,  55,  64,  140,  158, 

161,  445. 

Primary  phosphates,  269. 
Print,  photographic,  313. 
Problems,  21,  30,  49,  59,  86,  108,  132, 
146,  165,   180,  201,  224,  234, 
254,  264,  277,  283,  300,  318, 

330,  342,  353,  364,  372»  391, 
395,  436,  439,  440. 

Based  on  equations,  107. 
Producer  gas,  25. 
Products,  83. 

Addition,  204. 

Substitution,  203. 
Proof  spirit,  410. 
Propane,  203,  409. 
Properties  of  matter,  I,  2. 
Propyl,  409. 
Propylene,  202,  204. 
Proust,  77,  445. 
Prout,  398,  446. 
Prussian  blue,  388. 
Prussiate  of  potash,  red,  387. 

Yellow,  198,  387. 
Prussic  acid,  198. 
Puddling,  379. 
Pulp,  paper,  429. 
Purification,  water,  39. 
Purifiers,  gas,  212. 
Purple  of  Cassius,  317. 
Putty,  323. 
Pyrite,  386. 
Pyrogallic  acid,  431. 
Pyroligneous  acid,  415. 


582 


Index. 


Pyrolusite,  369. 
Pyromorphite,  357. 
Pyrophosphates,  269. 
Pyrophosphoric  acid,  269. 
Pyrosulphuric  acid,  252. 
Pyrrhotite,  373. 

Quadrivalent  elements,  176. 
Qualitative  analysis,  50,  242. 
Quantitative  analysis,  50. 
Quantitative  significance  of  equations, 

104. 

Quantivalence,  176. 
Quartation,  315. 
Quartz,  255,  256. 
Quartzite,  256. 
Quicklime,  324. 
Quicksilver,  338. 
Quinine,  433. 
Quinquivalent  elements,  176. 

Radical,  89,  150,  198. 

Organic,  406. 

Valence,  177. 
Rain  water,  37. 
Ramsay,  68,  69,  446. 
Rational  formula,  407. 
Rayleigh  and  Ramsay,  68. 
Reaction,  3. 

Acid,  90. 

Alkaline,  92. 

Chemical,  83. 

Illustrating  equation,  106. 

Neutral,  94. 
Realgar,  272,  273. 
Red  fire,  329. 

Hematite,  374. 

Paint,  340,  361,  384. 

Lead,  360. 

Liquor,  350,  417. 
Reduction,  15,  28,  55,  357. 

Process  for  lead,  358. 
Reducing  agent,  28. 

Flame,  222. 
Reference  books,  450. 


Refining  petroleum,  208. 
Relative  humidity,  66. 
Respiration,  16,  191. 
Retorts,  coal,  210. 
Reverberatory  furnace,  281,  282. 
Reversion,  271. 
Rhigolene,  208. 
Rhodium,  392. 
Rhodocroisite,  369. 
Rinmann's  green,  390. 
River  water,  38. 
Rochelle  powder,  290. 
Rock,  crystal,  255. 

Phosphate,  271. 
Rocks,  258. 

Decayed,  265. 

Phosphorus  from,  265. 

Silicates,  255. 
Roll  sulphur,  238. 
Rosaniline,  431. 
Rosendale  cement,  325. 
Rose's  metal,  275. 
Rouge,  384. 
Royal  water,  1 60. 
Rubidium,  284,  299,  413. 
Ruby,  347. 

Ore,  301. 
Rum,  411. 

Run,  water  gas,  213. 
Rusting  of  iron,  383. 
Rutherford,  63,  64,  446. 
Ruthenium,  392. 

Saccharose,  423. 
Safety  lamp,  221. 
Sal  ammoniac,  151. 
Saleratus,  290. 
Salicylic  acid,  431. 
Sal  soda,  289. 
Salt,  94. 

Acid,  96. 

As  glaze,  352. 

Basic,  96. 

Cake,  288. 

Common,  286, 


Index. 


583 


Salt,  continued. 

From  White  Sea,  287. 

Glauber's,  292. 

Microcosmic,  269. 

Preparation  of  common,  287. 

Springs,  228. 
Saltpeter,  Chili,  231,  292. 

Source,  295. 
Salts,  94. 

.  Action  on  litmus,  94. 

Ammonium,  150. 

And  ionization,  129. 

Epsom,  333. 

Ethereal,  419. 

Formation,  94. 

General  properties,  88. 

Haloid,  225. 

In  ocean,  38,  287. 

Nomenclature,  95. 

Normal,  96. 

Organic,  419. 

Smelling,  152. 
Sand,  255,  256. 

And  hydrofluoric  acid,  228. 

Blast,  257. 
Sandstone,  256. 
Saponification,  422. 
Sapphire,  347. 
Satin  spar,  326. 
Saturated  compounds,  177. 

Hydrocarbons,  203. 

Point  of  air,  66. 

Solution,  44. 
Scandium,  401. 
Scheele,  16,  18,  64,  133,  265, 

446. 

Scheele's  green,  273. 
Scrubber,  212. 
Seal,  210. 
Sea  water,  salts  in,  287. 

Silver  in,  308. 

Secondary  phosphates,  269. 
Seidlitz  powders,  290,  418. 
Selenite,  326. 
Selenium,  252. 


Series,  homologous,  202. 

Paraffin,  203. 
Serpentine,  331. 
Shell,  in  limestone,  322. 

Rock,  322. 
Shot,  360. 

Sicily,  sulphur  from,  236. 
Siderite,  373,  387. 

Siemens-Martin  process  for  steel,  382. 
Silica,  255. 

And  plants,  257. 

Deposition,  258. 

From  springs,  258. 

Hydrated,  256. 

Soluble,  258. 
Silicates,  257,  258. 
Siliceous  sinter,  258. 
Silicic  acid,  257. 
Silicides,  258. 

Carbon,  117. 
Silicified  wood,  256,  257. 
Silicon,  255. 

Bronze,  305. 

Carbide,  117. 

Tetrafluoride,  228,  257. 
Silicon  dioxide,  255. 

Properties,  256. 

Varieties,  255. 
Silver,  308. 

Acetate,  molecular  weight,  1 70. 

Alloys,  308,  311. 

Amalgam,  309. 

Amalgamation  process,  309. 

Bearing  lead,  308. 

Brick,  310. 

Bromide,  312. 

Chloride,  308,  309,  312. 

Coins,  311. 

Compounds,  312. 

Compounds  and  light,  312,  313. 

Determination   of  atomic  weight, 
171. 

Distribution,  309. 

German,  306. 

Glance,  308. 


584 


Index. 


Silver,  continued. 

Halogens,  solubility,  252. 

History,  308. 

Horn,  133,  308. 

In  sea  water,  308. 

Iodide,  312. 

Metallurgy,  309,  310. 

Name,  308. 

Nitrate,  312. 

Ores,  308. 

Oxidized,  311. 

Plating,  311,  312. 

Production,  308. 

Properties,  310. 

Pure,  310. 

Separation  from  gold,  315. 

Specific  heat,  173. 

Sterling,  311. 

Sulphides,  308,  311. 

Tarnishing,  311. 

Test,  312. 

Water,  338. 

World's  supply,  308. 
Silverware,  blackening,  242,  311, 
Simplest  formula,  104,  175. 
Sinter,  siliceous,  258. 
Sirius,  23, 
Sirup,  table,  426. 
Slag,  281,  324,  375. 

Phosphate,  271. 
Slaked  lime,  324. 
Slate,  343. 
Smalt,  390. 
Smelling  salts,  152. 
Smelting,  281.     See  Metallurgy. 
Smithsonite,  334. 
Smokeless  gunpowder,  428. 
Snow  crystals,  35. 
Soap,  420,  422. 

And  hard  water,  327. 

Boiling  process,  423. 

Cold  process,  423. 

Hard,  422. 

Soft,  422. 

White,  422. 


Soap,  continued. 

Yellow,  423. 
Soapstone,  331. 
Soda,  289,  290. 

Ash,  289. 

Baking,  290. 

Cooking,  290. 

Crystals,  289. 

Washing,  289. 

Water,  42,  90,  193. 
Sodium,  284. 

Acetate,  417. 

Alum,  349. 

Aluminate,  348,  349. 

Amalgam,  292,  339. 

And  water,  24,  51. 

Arsenate,  273. 

Arsenite,  273. 

Bicarbonate,  195,  289. 

Carbonate,  284,  288,  289. 

Chloride,  286,  287. 

Cyanide,  286,  293. 

Dioxide,  293. 

Discovery,  284. 

Hydroxide,  290,  291,  292. 

Hypochlorite,  139. 

Hyposulphite,  138,  252. 

lodate,  230. 

Iodide,  319. 

Lactate,  418. 

Manganate,  372. 

Manufacture,  284,  285. 

Monoxide,  293. 

Name,  284. 

Nitrate,  292,  293. 

Oxides,  286. 

Peroxide,  286,  293. 

Preservation,  286. 

Properties,  285. 

Silicate,  258. 

Stannate,  357. 

Sulphate,  292. 

Sulphide,  288. 

Sulphite,  243. 

Test,  141,  286. 


Index. 


585 


Sodium,  continued. 

Thiosulphate,  252. 

Tungstate,  369. 

Uses,  286. 
Soft  coal,  185,  189. 

Water,  37,  327. 
Solder,  356,  360. 
Soldering,  263. 
Solid  carbon  dioxide,  193. 
Solids,  solution,  43. 

Table,  44. 
Soluble  glass,  257. 

Silica,  258. 

Sulphate,  test,  251. 
Solute,  41. 
Solution,  41,  126. 

And  chemical  action,  47. 

And  electrolysis,  126. 

Boiling  point,  127. 

Freezing  point,  127,  128. 

Gases,  41. 

Labarraque's,  139. 

Liquids,  43. 

Nature,  48. 

Saturated,  44. 

Solids,  43. 

Supersaturated,  45. 

Terms,  41. 

Thermal  phenomena,  47. 
Solvay  process  for  sodium  carbonate, 

289. 
Solvent,  41. 

Universal,  43. 
Souring,  138. 

Sour  milk  in  cooking,  418. 
Specific  gravity  of  metals,  279. 
Specific  heat,  172. 

Law,  172. 

Table,  173. 
Spectra,  402. 

Nebulae,  404. 

Stars,  404. 
Spectroscope,  23,  402. 

Discovery  by,  284,  404. 
Spectrum,  401. 


Spectrum,  continued. 

Absorptive,  403. 

Analysis,  401,  403. 

Banded,  402. 

Bright  line,  402. 

Dark  line,  402. 

Sunlight,  403. 
Speculum  metal,  306. 
Spelter,  335. 
Sperrylite,  392. 
Sphalerite,  334. 
Spiegel  iron,  369,  378. 
Spinel,  ruby,  347. 
Spinels,  347. 
Spirit  of  salt,  140. 
Spirits,  hartshorn,  147. 
Spongy  platinum,  392,  393. 
Springs,  mineral,  37,  42. 
Stable  refuse,  271. 
Stack,  210. 
Stahl,  1 6,  446. 
Stalactite,  321. 
Stalagmite,  321. 
Stamp,  mill,  280. 
Standard  conditions,  19. 

Wax  candle,  216. 
Stannic  chloride,  357. 

Oxide,  356. 

Stannous  chloride,  356. 
Stannum,  354. 
Starch,  426,  427. 

And  potassium  salts,  298. 

Test,  232,  427. 
Stas,  171,  398,  446. 
Stassfurt  deposits,   133,  228,  261,  293, 

331-  . 
Steam,  36. 
Stearic  acid,  417. 
Stearin,  420. 

Candles,  422. 
Steel,  and  coke,  190. 

Bessemer,  381. 

Chrome,  366. 

Crucible,  381. 

Harveyized,  380. 


586 


Index. 


Steel,  and  coke,  continued. 

Manufacture,  380. 

Nickel,  389. 

Open  hearth,  383. 

Properties,  380. 

Tempering,  380. 

Uses,  383. 
Sterling  silver,  311. 
Stibine,  274. 
Stibium,  274. 
Stibnite,  274. 
Still,  40,  379. 
Stone,  artificial,  258. 

Ice,  350. 
Stoneware,  352. 
Stove  polish,  183. 
Strass,  260. 
Stream  tin,  356. 

Striking  back,  Bunsen  flame,  220. 
Strontia,  328. 
Strontium,  328. 

Carbonate,  328. 

Hydroxide,  328. 

Nitrate,  328. 

Oxide,  328. 

Sulphate,  328. 

Sulphide,  329. 

Test,  329. 

Structural  formulas,  178. 
Stucco,  327. 
Sublimate,  151. 

Corrosive,  340. 
Sublimation,  151,  440. 
Subnitrate  of  bismuth,  276. 
Substitution,  3,  203. 

Products,  203. 
Sucrose,  423. 
Suffix,  -ate,  95. 

-ic,  91,  144. 

-ide,  95. 

-ite,  95. 

-ous,  91,  144. 
Sugar,  423. 

Barley,  424. 

Beet,  424. 


Sugar,  continued. 

Brown,  424. 

Cane,  423,  424. 

Fermentation,  410. 

Fruit,  425.     . 

Granulated,  425. 

Grape,  425,  426. 

Kinds,  423. 

Of  lead,  363. 

Of  milk,  425. 

Raw,  424. 

Refining,  425. 

Term,  423. 

Test,  426. 

White,  424. 
Suint,  293. 

Potassium  carbonate  from,  297. 
Sulphates,  235,  251. 

Acid,  251. 

Important,  251. 

Normal,  251. 

Test,  141,  251. 
Sulphides,  238,  241. 

Color,  242. 

Native,  235. 

Solubility,  242. 
Sulphites,  245. 

Acid  calcium,  245. 

Acid  sodium,  245. 

Sodium,  243. 

Sulphur  dioxide  from,  243. 
Sulphur,  235. 

Action  with  heat,  238. 

Allotropic  modifications,  239. 

Amorphous,  239,  240. 

And  metals,  238. 

And  silver,  311. 

Burning,  245. 

Compounds,  240. 

Crystallized,  239. 

Dioxide,  242,  244,  245. 

Extraction,  236. 

Flowers,  238. 

Formation,  235. 

Forms,  239. 


Index. 


58? 


Sulphur,  continued. 

Free,  235. 

In  human  body,  236. 

In  United  States,  236. 

In  volcanic  districts,  235. 

Kiln,  236. 

Milk  of,  240. 

Monoclinic,  239. 

Native,  235. 

Orthorhombic,  239. 

Properties,  238. 

Purification,  237. 

Roll,  238. 

Source,  236. 

Springs,  37,  235. 

Trioxide,  245,  246. 

Use,  240,  252. 

Vapor  density,  238. 

Water,  37. 

Sulphuretted  hydrogen,  240. 
Sulphuric  acid,  246. 

And  organic  matter,  250. 

And  water,  250. 

Chemical  changes  in  making,  248. 

Concentration,  249. 

From  pyrites,  386. 

Fuming,  251. 

Impurities,  363. 

Manufacture,  246,  248,  249. 

Nordhausen,  252. 

Plant,  247-248. 

Properties,  250. 

Reduction,  250. 

Test,  251. 

Uses,  251. 

Sulphuric  ether,  414. 
Sulphurous  acid,  244,  245. 

Anhydride,  245. 
Sulphocyanic  acid,  198. 
Sun,  elements  in,  23,  404. 
Sunlight  and  carbon  dioxide,  194. 

Chemical  action,  m. 

Nitric  acid,  156. 
Superheater,  213. 
Superphosphate  of  lime,  271. 


Supersaturated  solution,  45. 
Supporter  of  combustion,  15. 
Sylvite,  294. 
Symbols,  81. 

And  atomic  weights,  103. 

Chemical,  8. 

Latin,  8. 

Table,  448,  449. 
Synthesis,  3,  50. 

Table  salt,  287. 

Tables,  atomic  weights,  448,  449. 

Borax  bead  colors,  262. 

Composition  of  coal,  186. 

Composition  of  natural  waters,  38. 

Equivalents,  100. 

Famous  chemists,  447. 

Important  elements,  6. 

lonization,  127. 

Latin  symbols,  8. 

Metric  equivalents,  438. 

Metric  system,  437. 

Metric  transformations,  438. 

Multiple  proportions,  78,  79. 

Periodic,  399. 

Solubility  of  carbon  dioxide,  42. 

Solubility  of  solids,  44. 

Specific  heats,  173. 

Uncommon  elements,  7. 

Water  in  food,  32. 
Talc,  331. 
Tallow,  421. 
Tannic  acid,  432. 
Tannin,  432. 
Tanning,  433. 
Tar,  213. 

Extractor,  212. 

Well,  210. 
Tartar,  crude,  418. 

Emetic,  274,  419. 
Tartaric  acid,  418. 
Tea,  433. 
Tellurides,  314. 
Tellurium,  252. 

Compounds,  314. 


588 


Index. 


Temperature  and  luminosity,  218. 

Kindling,  113,  218. 

Low,  204. 

Standard,  19. 
Tempering,  380. 
Temporary  hardness,  327. 
Tension  of  water  vapor,  36. 
Terms,  electrochemical,  120. 
Terra  cotta,  352. 
Tests,  acetic  acid,  419. 

Alcohol,  419. 

Aluminium,  347. 

Antimony,  275. 

Arsenic,  273. 

Barium,  329. 

Bismuth,  276. 

Borax  bead,  262. 

Boron,  261. 

Cadmium,  337. 

Calcium,  328. 

Carbon,  189. 

Carbon  dioxide,  192,  325. 

Chloride,  144. 

Chromium,  367,  368. 

Cobalt,  390. 

Copper,  304. 

Gold!,  317. 

Hydrochloric  acid,  144. 

Hydrogen  sulphide,  242. 

Ions,  129. 

Iron,  388. 

Lead,  363. 

Lithium,  298. 

Manganese,  372. 

Marsh's,  for  arsenic,  273. 

Nickel,  389. 

Nitrates,  159. 

Nitric  acid,  159. 

Potassium,  294. 

Silver,  312. 

Sodium,  286. 

Starch,  427. 

Strontium,  329. 

Sugar,  426. 

Sulphate,  insoluble,  251. 


Tests,  continued. 

Sulphate,  soluble,  251. 

Sulphuric  acid,  251. 

Zinc,  337. 
Tetrads,  176. 
Tetragonal  crystals,  441. 
Theine,  433. 
Theory,  75. 

Atomic,  79. 

Electrolysis,  125. 

Electrolytic  dissociation,   125,  126. 
Thermal  equation,  112. 
Thermometers,  439. 
Thiosulphate,  sodium,  252. 
Thomas-Gilchrist    process    for    steel, 

382. 

Tiles,  352. 
Tin,  354. 

Alloys,  356. 

Amalgam,  339,  356. 

Block,  355. 

Crystals,  356. 

Dioxide,  354,  356. 

Foil,  356. 

History,  354. 

Interaction  with  metals,  355. 

Metallurgy,  354. 

Ore,  354. 

Oxymuriate,  357. 

Plate,  355. 

Production,  354,  356. 

Properties,  355. 

Purification,  355. 

Stone,  354. 

Stream,  356. 

Uses,  355. 
Tinkel,  261. 
Tinware,  355. 
Tobacco,  433. 
Toluene,  202,  430. 
Toluidine,  431. 
Toning,  in  photography,  313. 
Topaz,  347. 
Travertine,  322. 
Triacid  base,  94. 


Index. 


589 


Triads,  176. 
Tribasic  acid,  92. 
Triclinic  crystals,  442. 
Trivalent  elements,  176. 
Tungsten,  369. 
Turnbull's  blue,  388. 
Turquoise,  347. 
Tuscany,  boric  acid  from,  261. 
Tuyeres,  376. 
Type  metal,  360. 
Water,  89. 

Univalent  elements,  176. 
Unsaturated  compounds,  177. 

Hydrocarbons,  204,  206. 
Uranium,  369. 

Salts,  369. 

Specific  heat,  173. 
Urea,  405. 

Valence,  176. 

Classification  by,  397. 

Representation,  407. 
Valentine,  Basil,  246. 
Van  Helmont,  196,  446. 
Van't  Hoff,  446. 
Vapor  density,  169. 

And  molecular  weight,  168. 

Iodine,  232. 

Mercury,  339. 

Sulphur,  238. 

Zinc,  336. 
Vapor  tension,  36. 
Varec,  231. 
Vaseline,  209. 

Vegetable  matter  and  coal,  184-185. 
Vein  mining,  315. 
Venetian  red,  384. 
Verdigris,  417. 
Vermilion,  340. 
Vinegar,  90. 

Preparation,  415. 

Quick  process,  416. 

Wood,  415. 
Vital  force,  405. 


Vitriol,  blue,  307. 

Green,  385. 

Oil  of,  92,  246. 

White,  337. 

Volatile  alkali,  93,  149. 
Volta,  119. 
Voltaic  cell,  119. 
Volume  equation,  175. 
Volumetric,  53. 

Composition  of  air,  64. 

Composition  of  water,  53,  55,  57. 

Washing  soda,  289. 

Washington  monument,  cap,  345. 

Water,  31. 

Analysis,  39. 

And  chlorine,  51. 

And  hydrogen,  50. 

And  oxygen,  51. 

And  sodium,  24,  51. 

As  solvent,  32,  33. 

Baryta,  329. 

Boiling  point,  36,  439. 

Chalybeate,  37,  387. 

Chlorine,  135. 

Composition,  25,  27. 

Density,  34. 

Distilled,  40. 

Drinking,  39. 

Electrolysis,  52,  123. 

Expansion,  34. 

Freezing,  34,  439. 

From  burning  hydrogen,  27. 

Function  in  nature,  32. 

Gas,  25,  196,  213,  214,  215. 

Glass,  25^. 

Gravimetric  composition,  55,  57. 

Hard,  37,  327. 

Hardness,  327. 

Hydrogen  sulphide,  241. 

Industrial  application,  33. 

In  food,  31,  32. 

In  human  body,  32. 

In  liquid  state,  31. 

In  vegetables,  31,  32. 


590 


Index. 


Water,  continued. 

Javelle's,  139. 

Lithia,  298. 

Mineral,  37. 

Natural,  37. 

Occurrence  in  nature,  31. 

Ocean,  38. 

Of  crystallization,  45,  46. 

Physical  properties  of  pure,  33. 

Purification,  39,  371. 

Quantitative  composition,  53. 

Rain,  37. 

River,  38. 

Silver,  338. 

Soda,  42. 

Soft.  37,  327. 

Type,  89. 

Underground,  37. 

Volumetric  composition,  53,  55,  57. 
Water  vapor,  31,  36. 

Condensed,  31,  36. 

In  atmosphere,  62,  66. 
Watt,  55. 

Wax,  paraffin,  209. 
Welding  iron,  379. 
Weldon,  mud,  370. 

Process,  134,  370. 
Welsbach  light,  222. 
WTet  process,  47,  282. 
Whetstone,  256. 
Whisky,  411. 
White  arsenic,  272. 

Cast  iron,  378. 

Lead,  361. 

Magnesia,  370. 

Metal,  306. 

Paint,  242,  336,  362. 

Vitriol,  337. 
Whitewash,  326. 
Whiting,  323. 
Willemite,  334. 
Willson,  1 1 6. 
Winds,  62. 
Wine,  in,  297. 
Witherite,  329. 


Wohler,  343,  405,  446. 
Wood  alcohol,  409. 

Ashes,  297. 

Charcoal,  187. 

Petrified,  256,  257,  258. 

Preserving,  337. 

Silicified,  256,  257. 

Spirit,  409. 

Vinegar,  415. 
Wood's  metal,  275,  337. 
Worm,  condenser,  40. 
Wrought  iron,  378. 

Xenon,  69,  404. 

Yeast,  410. 

In  bread-making,  427. 
Yellow  paint,  367. 

Zinc,  334. 

Alloys,  306,  336. 

Blende,  334. 

Carbonate,  334. 

Chloride,  122,  337. 

Deposits,  334. 

Determination  of  atomic  weight 

173- 

Dust,  335,  336. 

Hydroxide,  337. 

Metallurgy,  334. 

Ores,  334. 

Oxide,  334,  335,  336,  362. 

Production,  334. 

Properties,  335. 

Silicate,  334. 

Smelting,  334. 

Sulphate,  336. 

Sulphide,  334,  336. 

Test,  337. 

Uses,  336. 

Vapor  density,  336. 

White,  336. 
Zincates,  335,  337. 
Zincite,  334. 
Zero,  absolute,  439. 


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